SOLI 


01 


3.9 


Br 


NO, 


C10S 


Bra, 


I03 


OH 


SO, 


CrO4 


C204 


CO, 


4.G 


137.5 


92.56 
12.4 


30.34 
2.6 


6.6 
0.52 


0.38 


0.35 


142.9 


11.11 
0.62 


63.1 


30.27 
1.6 


108.0 
5.9 


m^Fz£g& 

L1SRIS 


COLLEGE  OF  AGRICULTURE 
DAVIS,  CALIFORNIA 


AT   18°, 


Zn 


Pb 


16.83 
1.15 


61.21 
3.30 


3.34 
0.24 


1.8 


5.0 


35.64 
2.8 


111.6 
6.5 


7.22 


1.3 
0.17 


0.001 


0.55 
0.020 


0.0025 
0.0315 


0.0035 
0.032 


0.031 


1.76 


4.74 
0.09 


0.006 
0.031 


1.48 
0.030 


4.95 
0.10 


0.22 


0.0323 
0.0410 


0.0,38 
0.0415 


0.0080 
0.0338 


0.0023 
0.0311 


0.063 


0.011 
0.036 


0.12 
0.006 


0.0046 
0.0.,26 


0.0011 
0.047 


0.02 


0.20 
0.015 


0.4 
0.03 


0.0356 
0.0443 


0.0013 
0.0313 


...JOl 
0.032 


35.43 

2.8 


73.0 
4.3 


0.03 
0.0027 


0.1 
0.01 


203.9 
9.2  0.05 


478.2 
9.8 


419 


0.006 
0.035 


117.8 


183.9 
5.3 


58.43 
1.8 


0.83 
0.02 


0.035 
0.0t5 


53.12 
3.1 


0.02 


0.08 
0.0,2 


0.07 
0.003 


51.66 
1.4 


150.6 
3.16 


1.3 
0.03 


0.043 


0.01 
0.0.,4 


0.0041 
0.0S13 


0.042 
0.065 


0.036 
0.044 


0.004? 
0.033? 


0.0,15 
0.055 


0.0,1 
0.043 


The  'tipper  number  in  each  square  gives  the  number  of  grams  of  the 
anhydrous  salt  held  in  solution  by  100  c.c.  of  water.  The  loiver  number  is 
the  molar  solubility,  i.e.,  the  number  of  moles  contained  in  one  liter  of  the 
saturated  solution.  For  some  other  solubilities,  see  pages  157  and  544. 


t^iA^U^  ' 


tbe  same  autbor 


A  Companion  Laboratory  Guide  to 
this  Text 

A  LABORATORY  OUTLINE  OF 
GENERAL    CHEMISTRY 

NEW    EDITION,   19O8 

Fourth  Edition,  Revised 

In  Collaboration    with   Dr.  W.  J.   HALE 

of  the   University   of  Michigan 
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INTRODUCTION 

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INORGANIC  CHEMISTRY 


BY 


ALEXANDER  SMITH 

PROFESSOR  OF  CHEMISTRY,  AND  HEAD  OF  THE  DEPARTMENT, 
COLUMBIA  UNIVERSITY 


1Rex>iset>  Edition 

THIRTY-FIRST    THOUSAND 


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1912 

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LIB 

COLLEGE  OF  AGRICULTURE 
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COPYRIGHT,   1906,   1906, 
BY 

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FIRST  EDITION,  MARCH  l,  1906. 
REPRINTED  JULY,  1906;  OCTOBER,  1906; 

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REPRINTED  APRIL,  1912. 


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BOSTON,     U.  8.  A. 


PREFACE 


THIS  book,  the  first  draft  of  which  was  written  six  years  ago,  is  the 
outgrowth  of  the  introductory  course  in  chemistry  which  the  author 
has  given  for  the  past  fifteen  years.  A  subject  undergoing  the  per- 
sistent, though  unconscious  criticism  of  keen  minds  should  gain  in 
self-consistency  and  coherence  as  it  is  presented  year  after  year.  For 
example,  an  answer  must  be  found  for  the  common  question,  "  Why 
does  the  chemistry  of  the  laboratory  differ  from  the  chemistry  of  the 
text-book  and  the  lecture  to  such  an  extent  that  they  seem  to  be  different 
sciences  ?  "  The  chemistry  of  the  laboratory  is,  of  course,  the  only 
real  chemistry,  and  that  of  the  lecture  must  be  somewhere  at  fault. 
The  student  neither  sees  nor  weighs  atoms,  for  instance,  and  so  the 
details  of  the  laboratory  experiment,  which  are  seen  and  studied,  become 
the  basis  of  the  whole  treatment.  The  atom  and  the  ion  assume  the 
role  of  merely  figurative  aids  in  the  description  of  the  facts.  Gradually 
the  conception  of  chemical  equilibrium  comes  to  contribute  the  major 
part  of  the  explanation  which  is  essential  to  the  evolution  of  a  system 
of  chemistry  founded  upon  experiment. 

In  the  choice  and  arrangement  of  the  material,  several  principles 
have  served  as  guides  : 

The  book  is  intended  primarily  for  students  beginning  the  study  of 
chemistry  in  a  college,  university,  or  professional  school.  It  is  assumed 
that  use  of  the  book  goes  hand  in  hand  with  systematically  arranged 
laboratory  work  in  general  chemistry.  The  first  four  chapters,  for  ex- 
ample, contain  a  discussion  of  a  few  typical  experiments.  They  appeal 
directly  to  experience  derived  from  the  performance  and  observation 
of  these  and  other  similar  experiments  in  the  laboratory  and  in  the 
class-room.  In  these  chapters  some  of  the  features  which  are  charac- 
teristic of  every  chemical  phenomenon  are  sought  out,  put  into  words, 
and  illustrated. 


+8111 


yi  PREFACE 

No  conception  is  denned,  and  no  generalization  or  law  is  developed, 
until  such  a  point  has  been  reached  that  applications  of  the  conception 
and  experimental  illustrations,  later  to  be  related  in  the  law,  have 
already  been  encountered,  and  there  is  about  to  be  occasion  for  further 
applications  and  illustrations  of  the  same  things  in  the  chapters 
immediately  succeeding.  In  these  chapters  the  applications  are 
frequent  and  explicit.  Later,  page  references  in  parentheses  con- 
tinue to  indicate  the  recurrence  of  examples  which  might  other- 
wise fail  to  be  noticed.  It  is  one  thing  to  come  to  know  a  principle 
of  the  science,  and  quite  another  thing  to  have  acquired,  by  constant 
repetition  of  the  process,  a  confirmed  habit  of  using  the  prin- 
ciple on  every  appropriate  occasion.  To  assist  still  further  in  the 
attainment  of  this  end,  an  attempt  is  made  in  the  last  six  chapters 
again  to  mention  and  illustrate,  by  way  of  review,  the  most  important 
principles  of  the  science. 

No  conception  or  principle  is  given  at  all,  unless,  in  its  most  elemen- 
tary aspects,  it  can  be  made  clear  to  a  beginner  ;  and  unless  it  is 
capable  of  numerous  applications  in  elementary  work;  and,  finally, 
unless  a  knowledge  of  it  is  of  material  use  in  organizing  and  unifying 
the  result  of  such  elementary  work. 

An  attempt  has  been  made  to  state  the  laws  and  to  define  the  con- 
ceptions of  the  science  in  terms  of  experimental  facts.  The  figurative 
language  of  hypothesis  has  been  employed  only  in  explanations. 

Familiarity  with  physical  conceptions  and  facts  is  so  indispensable 
to  the  chemist  that  no  apology  is  needed  for  the  rather  full  treatment 
which  some  of  them  have  received. 

No  two  chemists  would  agree  perfectly  in  regard  to  the  apportion- 
ment of  space.  The  processes  of  chemical  industry,  and  the  every-day 
applications  of  chemical  science,  cannot  all  be  mentioned.  These 
fields,  and  that  of  mineralogy,  can  be  represented  by  examples,  with- 
out the  incompleteness  of  the  result  being  in  any  way  a  detriment  to  a 
work  of  a  general  character.  Again,  a  dense  array  of  descriptive 
material,  unillumined  by  explanation,  is  a  positive  injury  to  an  intro- 
ductory treatise.  All  reference  to  historical  matters  cannot  be  omitted, 
but  a  logical  display  of  the  subject  can  be  achieved  with  comparatively 


PREFACE  vii 

little  of  the  history  Of  all  the  aspects  of  the  science,  the  theoretical 
is  thus  the  one  whose  treatment  is  susceptible  of  least  abbreviation. 

The  principles  of  chemical  equilibrium  are  (and  have  been  for  the 
past  half-century)  fully  as  much  required  for  intelligent  consideration 
of  the  simplest  experiment,  as  is  the  theory  of  combining  proportions 
itself.  Important  parts  of  the  theories  of  solutions  and  of  the  battery 
are  much  more  recent,  but  each  is  equally  indispensable  to  the  under- 
standing of  matters  which  cannot  long  be  withheld  from  the  notice  of 
the  beginner.  Surely  space  ought  not  to  be  saved  by  entire  omission  of 
essential  parts  of  the  chief  thing  that  makes  chemistry  at  all  worthy 
of  a  place  amongst  the  sciences.  Nor  may  we  attain  brevity,  no  matter 
how  great  the  temptation,  by  condensing  the  passages  on  theory  until 
they  reach  the  limit  of  comprehensibility  by  an  expert.  Without  clear 
exposition,  full  illustration,  and  frequent  application,  laws  and  princi- 
ples simply  repel,  or  worse  still,  mislead  the  beginner. 

We  reach  the  same  conclusion-  from  another  view-point.  Every 
student  should  have  access  to,  and  should  use,  reference  books  devoted 
especially  to  descriptive,  industrial,  historical,  and  physical  chemistry, 
and  to  mineralogy  and  crystallography  —  at  least  one  good  book  in 
each  of  these  five  subjects.  With  the  help  of  the  index,  the  veriest  tyro 
can  find  in  a  few  moments,  almost  anything  he  wants  in  four  out  of 
five  of  these  branches.  But  just  the  opposite  is  the  case  with  the 
theory.  Only  an  expert  realizes  what  information  he  is  in  need  of, 
and  knows  under  what  titles  to  look  for  it.  And  often  even  the 
expert  would  fail  to  understand  the  isolated  sentence  or  paragraph 
when  found.  In  a  large  proportion  of  connections  the  beginner  sim- 
ply cannot  use  such  a  book  for  rapid  reference  at  all.  In  many  lines, 
therefore,  much  may  be  left  to  outside  reading,  but  for  theory  almost 
no  dependence  can  be  placed  on  reference  work  in  other  books. 

For  the  reasons  enumerated  above,  an  unusually  large  proportion 
of  space  has  been  given  to  theoretical  matters.  The  actual  amount  of 
theory  is  no  greater  than  is  usual  in  books  of  the  same  class,  but  the 
explanations  are  often  fuller.  Even  so,  the  beginner  will  probably 
find  that  some  parts  form  reading  as  stiff  as  any  he  is  accustomed  to 
undertake,  without  complaint,  in  physics  or  mathematics.  It  can  only 


Vlll  PREFACE 

be  said  that  easily  read  modes  of  presenting  the  science  of  chemistry- 
are  apt  to  delude  the  beginner  into  thinking  he  has  mastered  the  sub- 
ject, when  in  reality  he  has  simply  been  steered  clear  of  the  chief  diffi- 
culties. 

The  order  of  topics  was  determined  by  many  considerations, 
jointly.  For  example,  in  the  first  week  of  his  work,  a  student  may 
encounter  experiments,  in  connection  with  which,  almost  every  part  of 
chemical  theory  might  usefully  be  discussed.  But  mastery  of  the 
theory  must  necessarily  come  bit  by  bit,  and  the  theory  is  therefore 
distributed  through  the  book.  Instead  of  being  introduced  as  soon  as 
a  fragment  offers  a  chance  for  explanation,  the  treatment  of  each  of  the 
various  theoretical  subjects,  as  far  as  possible,  has  been  postponed  until 
a  whole  chapter  could  be  devoted  to  it.  The  result  makes  subsequent 
reference  easier,  and  facilitates  alterations  in  the  order  of  study. 
Thus,  the  hypothesis  of  ions  is  not  mentioned  as  soon  as  it  well  might 
be,  because  satisfactory  treatment  of  it  must  follow  the  molecular  and 
atomic  hypotheses  of  which  it  is  an  extension,  and  because  the  full 
explanation  of  this  hypothesis  must  be  preceded  by  some  account  of  the 
phenomena  of  electrolysis  and  of  the  essential  properties  of  solutions, 
and,  also,  by  a  discussion  of  chemical  equilibrium,  a  subject  which  of 
necessity  presupposes  two  or  three  months'  work  iii  chemistry.  There 
is  another  disadvantage  which  arises  from  a  premature  explanation  of 
the  hypothesis  of  ionization.  When  it  appears  at  an  early  stage,  too 
long  an  interval  separates  this  subject  from  the  study  of  the  metallic 
elements,  and  the  details  are  largely  forgotten  before  the  field  for  their 
chief  employment  is  reached. 

The  paragraphs  in  smaller  type  are  not  intended  for  beginners,  but 
for  advanced  students  and  teachers,  which  accounts  for  the  fact  that 
reference  will  frequently  be  found  in  them  to  subjects  treated  system- 
atically in  later  chapters. 

The  exercises  and  problems  are  simply  samples  of  some  of  the 
various  kinds  of  questions  which  might  be  raised  in  dozens  at  the  end 
of  every  chapter. 

Recent  works  on  general  chemistry  have  been  consulted  during  the 
revision  of  the  manuscript.  Of  these  A.  A.  Noyes'  admirable  General 


PREFACE  IX 


Principles,  Ostwald's  Grundlinien  —  a  veritable  tour  de  force,  and  Blox- 
am's  Chemistry  may  be  mentioned  as  having  proved  most  suggestive. 
The  author  owes  special  thanks  to  several  friends  who  have  undertaken 
the  toilsome  work  of  reading  part  or  all  of  the  book  in  manuscript  or  in 
proof,  and  in  particular  to  his  colleagues  Messrs.  Julius  Stieglitz,  H.  N. 
McCoy,  L.  W.  Jones,  and  E.  S.  Hall,  to  Pr.  J.  B.  Tingle  of  Johns 
Hopkins  University,  to  Mr.  C.  M.  Wirick  of  the  R.  T.  Crane  High 
School,  Chicago,  and  to  Mr.  Maurice  Pincoffs  of  Chicago.  The  author 
alone  is  responsible  for  any  defects  which  may  be  inherent  in  the  plan 
of  the  book  and  for  errors  which  may  have  escaped  detection,  but  must 
gratefully  acknowledge  the  very  great  benefit  the  book  has  derived 
from  the  friendly  criticism  of  these  gentlemen.  Other  corrections  or 
suggestions  will  be  gladly  received  by  the  author. 

CHICAGO,  January,  1906.  ALEXANDER  SMITH. 


CONTENTS* 


CHAPTERS  PAGES 

I.     Introductory  I 3 

The  Scientific  Method  —  Three  Illustrative  Chemical  Phenom- 
ena —  Two  Characteristics  of  Chemical  Phenomena. 

II.     Introductory  II 17 

Conservation  of  Mass  —  Physical  Concomitants  of  Chemical 
Change  —  Energy — Applications  of  the  Conception  of  Energy 
in  Chemistry  —  Free  Energy  —  Chemical  Activity  and  its 
Cause  —  Of  Simple  and  Compound  Substances  —  Some  of  the 
Fundamental  Ideas  Used  by  Chemists  and  the  Correspond- 
ing Terms  —  Methods  of  Work  and  Observation  in  Chem- 
istry. 

III.  Introductory  III 41 

The  Law  of  Definite  Proportions  —  The  Law  of  Multiple  Pro- 
portions—  The  Measurement  of  Combining  Weights  — 
The  Law  of  Combining  Weights  —  Equivalents  —  Atomic 
Weights. 

IV.  Introductory  IV 53 

Symbols,  Formulae,  and  Equations  —  Units  of  Measurement 
in  Chemical  Work  —  Calculations  in  Chemistry. 

V.     Oxygen 61 

Preparation  of  Simple  Substances  —  Preparation  of  Oxygen  — 
Physical  Properties  of  Oxygen  —  Specific  Chemical  Properties 
—  Chemical  Properties  of  Oxygen  —  The  Making  of  Equations 
Again — Oxides  —  Combustion — Oxidation — Means  of  Alter- 
ing the  Speed  of  a  Given  Chemical  Action :  By  Change  of 
Temperature  —  Rapid  Self-sustaining  Chemical  Action  and 
Means  of  Initiating  It  —  Other  Means  of  Altering  the  Speed 
of  a  Given  Chemical  Change  :  By  Change  in  Concentration  ; 
By  Catalysis;  By  Solution  —  Thermochemistry. 

VI.     The  Measurement  of  Quantity  in  Gases 80 

The  Variable  Concentration  of  Gases  —  The  Method  of  Allow- 
ing for  Varying  Concentration  in  Measuring  Quantity  in 
Gases  —  The  Relation  of  the  Volume  of  a  Gas  to  Temperature 
— Mixed  Gases  —  Densities  of  Gases  —  Vapor  Densities  of 
Liquids  and  Solids. 

*  The  titles  of  most  of  the  descriptive  paragraphs  have  been  omitted  from  this  table 
because  it  is  easier  to  find  matters  of  this  nature  by  consulting  the  index.  The  titles  of  all 
the  theoretical  paragraphs,  however,  have  been  included. 

XI 


xii  CONTENTS 

CHAPTERS  PAGES 

VII.     Hydrogen 92 

Acids  —  Preparation  of  Hydrogen  —  Tests  —  Displacement 

Valence  —  The  Valence  of  Radicals  —  How  to  Ascertain  the 
Valence  of  an  Element  or  Radical  —  Physical  Properties  of 
Hydrogen  —  Diffusion —  Chemical  Properties  of  Hydrogen  — 
A  False  Use  of  the  Word  "Affinity"  in  Explanation  of 
Chemical  Actions  —  The  Speed  of  Chemical  Actions  :  A 
Means  of  Measuring  Activity. 

VIII.     Water 113 

Natural  Waters  —  Physical  Properties  of  Water  —  Ice  — 
Steam  and  Aqueous  Tension  —  Chemical  Properties  of 
Water  —  Hydrates  —  Composition  of  Water  —  Gay-Lussac's 
Law  of  Combining  Volumes. 

IX.     The  Kinetic-Molecular  Hypothesis 128 

Kinetic-Molecular  Hypothesis  Applied  to  Gases  —  Critical 
Phenomena — Kinetic  Hypothesis  Applied  to  Liquids — Dif- 
fusion in  Liquids — Kinetic  Hypothesis  Applied  to  Solids  — 
Crystal  Forms  —  Crystal  Structure  —  Molecular  Magnitudes 

—  Formulative  and  Stochastic  Hypotheses. 

X.     Solution 146 

General  Properties  of  Solutions  —  The  Scope  of  the  Word  — 
Limits  of  Solubility  —  Recognition  and  Measurement  of  Solu- 
bility —  Terminology  —  Solution  one  of  the  Physical  States  of 
Aggregation  of  Matter  —  Kinetic-Molecular  Hypothesis  Ap- 
plied to  the  State  of  Solution  —  Kinetic-Molecular  Hypothesis 
Applied  to  the  Process  of  Solution  —  Independent  Solubility 

—  Solution  of  a  Gas  in  a  Liquid  —  Two  Immiscible  Solvents  : 
Law  of  Partition  —  Influence  of  Temperature  on  Solubility  — 
Phases  —  Equilibrium  in  a  Saturated  Solution — Metastable 
Condition  —  Saturation  —  Properties    of    Solutions   Propor- 
tional to  Concentration  :  Vapor  Tension  —  Freezing-Points 
of  Solutions  —  Densities  of  Solutions  —  Heat  of  Solution  — 
Definition  of  a  Solution  —  Application  in  Chemical  Work. 

XL    Chlorine  and  Hydrogen  Chloride 168 

Chlorine  —  Chemical  Relations  of  the  Element  —  Hydrogen 
Chloride  —  The  Kinetic  Hypothesis  Applied  to  the  Interaction 
of  Sulphuric  Acid  and  Salt  —  Classification  of  Chemical  In- 
teractions and  Exercises  Thereon  —  SUMMARY  or  PRINCIPLES. 

XII.     Molecular  Weights  and  Atomic  Weights  190 

Meaning  of  Avogadro's  Hypothesis. 

MOLECULAR  WEIGHTS  :  The  Relative  Weights  of  Molecules  — 
Molecular  Weights  —  The  Gram-Molecular  (Molar)  Volume. 

ATOMIC  WEIGHTS  :  Determination  of  the   Atomic  Weight   of 
Each  Element  —  Advantages  of  Atomic  Weights  over  Equiva- 


CONTENTS  xiii 

CHAPTERS  PAGES 

lents  —  Dulong  and  Petit's  Law,  an  Alternative  Means  of 
Determining  Atomic  Weights. 

MOLECULAR  FORMULAE:  Molecular  Formulae  of  Compounds  — 
The  Molecular  Weights  and  Formulse  of  Elementary  Sub- 
stances —  Further  Discussion  of  the  Molecular  Formulae  of 
Elementary  Substances. 

APPLICATIONS:  Interactions  Between  Gases  —  Molecular  Equa- 
tions —  Arithmetical  Problems  —  Cases  of  Dissociation  — 
Finding  the  Atomic  Weight  of  a  New  Element  —  Replies  to 
Questions  and  Difficulties  —  Exercises. 

XIII.  The  Atomic  Hypothesis 217 

XIV.  The  Halogen  Family 226 

The  Chemical  Relations  of  Elements  —  The  Chemical  Rela- 
tions of  the  Halogens  —  Bromine  —  A  Plan  for  Making  Very 
Complex  Equations  —  Hydrogen  Bromide  — •  Iodine — Hydro- 
gen Iodide  —  Fluorine  —  Hydrogen  Fluoride  —  Association — 
The  Halogens  as  a  Family  —  Compounds  of  the  Halogens 
with  Each  Other. 

XV.      Chemical  Equilibrium 246 

Reversible  Actions  —  Kinetic  Explanation  —  The  Influence 
of  Homogeneous  Mixture  —  The  Influence  of  Molecular  Con- 
centration — Law  Connecting  Molar  Concentration  and  Speed 
of  Reaction— The  Condition  for  Chemical  Equilibrium — The 
Effect  of  Changes  of  Volume  on  Chemical  Equilibrium 
—  Heterogeneous  Equilibrium  —  Applications  in  Chemistry  : 
Displacement  of  Equilibria  by  Changes  Affecting  Concentra- 
tion —  Affinity  vs.  Solubility  —  Displacement  of  Equilibria  by 
Changes  in  Temperature  :  van  »t  HofTs  Law  and  Le  Chate- 
lier's  Law — SUMMARY  OP  PRINCIPLES. 

XVI.    Oxides  and  Oxygen  Acids  of  the  Halogens 263 

Compounds  of  Chlorine  Containing  Oxygen  —  Nomenclature 
of  Acids  and  Salts  —  Salts  and  Double  Decomposition  — 
Hypochlorites  and  Hypochlorous  Acid  —  Hypochlorous 
Anhydride  —  Hypochlorous  Acid  as  an  Oxidizing  Agent : 
Bleaching  —  Thermochemistry  of  Hypochlorous  Acid  — 
Simultaneous,  Independent  Chemical  Changes  in  the  Same 
Substance  —  Chlorates  —  The  Separation  of  Substances  by 
their  Solubility  —  Chloric  Acid  —  Making  of  Equations 
Again  —  Chlorine  Dioxide  :  Chlorous  Acid  —  Perchlorates, 
Perchloric  Acid,  and  Perchloric  Anhydride  —  Oxygen  Acids 
of  Bromine  —  Oxides  and  Oxygen  Acids  of  Iodine  — lodates 
and  lodic  Acid  —  Various  Acids  Derived  from  One  Anhy- 
dride—  Periodates  and  Periodic  Acid  —  Chemical  Rela- 
tions. 


XIV 


CONTENTS 


CHAPTERS  PAGES 

XVII.  Dissociation  in  Solution  281 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution. 
OSMOTIC  PRESSURE  :  Phenomena  Produced  by  Osmotic  Pressure 

—  The  Phenomena  a  Logical  Consequence  of  Semi -permea- 
bility—  Measurement  of  Osmotic  Pressure — Osmotic  Pres- 
sure and  Concentration  —  Osmotic  Pressure  and  Temperature 

—  An  Analogue  of  Avogadro's  Hypothesis  —  Determination 
of  Molecular  Weights  —  Osmotic  Pressure  and  Dissociation 
in  Solution. 

DEPRESSION  IN  THE  FREEZING-POINT  OF  A  SOLVENT  :  Measure- 
ment of  Freezing-Points  —  Laws  of  Freezing-point  Depres- 
sion—  Determination  of  Molecular  Weights  —  Freezing- 
Points  and  Dissociation  in  Solution  —  Boiling-Points  and 
Dissociation  in  Solution  —  Comparison  of  the  Results  of 
the  Three  Methods. 

THE  APPLICATION  OP  THESE  CONCLUSIONS  IN  CHEMISTRY  : 
The  Constitution  of  Solutions  of  Acids,  Bases,  and  Salts  — 
Nomenclature  :  The  Ionic  Hypothesis  — Ionic  Equilibrium  — 
Appendix  :  Recent  Measurements  of  Osmotic  Pressure. 

XVIII.  Ozone  and  Hydrogen  Peroxide 300 

Ozone  —  Hydrogen  Peroxide  —  The  Interaction  of  Barium 
Peroxide  and  Sulphuric  Acid  —  Thermochemistry  of  Hydro- 
gen Peroxide  —  Peroxides  :  Chemical  Constitution  and 
Molecular  Structure. 

XIX.    Electrolysis 310 

Introductory  —  Chemical  Changes  Connected  with  Electro- 
lysis —  Ionic  Migration  —  Relative  Speed  of  Migration  of 
Different  Ions  —  Faraday's  Laws  —  The  Ionic  Hypothesis  — 
Difficulties  Presented  by  this  Hypothesis  —  Amounts  of 
Electricity  on  the  Ions  —  Re'surne'  and  Nomenclature  — 
Actual  Quantities  of  Electricity  Concerned  in  Electrolysis 
— The  Electrical  Energy  Required  to  Decompose  Different 
Compounds  —  Polarization  —  Conductivity  for  Electricity  — 
Degrees  of  lonization  of  Common  Substances  :  Acids,  Bases, 
Salts  —  Degree  of  lonization  of  Water  —  General  Remarks 
on  these  Values  —  Comparison  with  the  Results  Obtained 
by  Other  Methods. 

XX.     The  Chemical  Behavior  of  Ionic  Substances 334 

Solutions  of  lonogens  are  Mixtures  —  Each  Kind  of  Ion  in  a 
Mixture  Acts  Independently. 

SALTS,  IONIC  DOUBLE  DECOMPOSITION,  PRECIPITATION  :  Salts 
— Double  Decomposition  of  Salts  in  Solution  —  Solution 
and  Precipitation  of  Salts  —  Individual,  Specific  Chemical 
Properties  of  Each  Ionic  Material  —  Application  in  Chemical 
Analysis  —  Hydrolysis  of  Salts. 


CONTENTS  XV 

CHAPTERS  PAOKS 

ACIDS  AND  BASES  AND  THEIR  DOUBLE  DECOMPOSITION  WITH 
SALTS  :  Hydrogen  Salts  —  Hydrion  — Modes  of  lonization  of 
Acids  —  Activity  of  Acids  —  Salts  of  Hydroxyl  —  Hydrox- 

idion Ionic  Double  Decomposition  and  Precipitation  of 

Acids  and  Bases — Ionic  Double  Decomposition  and  Activity. 
NEUTRALIZATION:   Acidimetry  and  Alkalimetry  —  Theory  of 
Neutralization  —  Indicators  —  Neutralization  of  Little-ion- 
ized   Substances  —  Thermochemistry   of    Neutralization  — 
Volume  Change  in   Neutralization. 
MIXED  IONOGENS  AND  DOUBLE  SALTS  :  Acid  Salts  — Basic  Salts 

—  Mixed  Salts  —  Double  Salts. 

KINDS  OF  IONIC  CHEMICAL  CHANGES  :  Disunion  and  Combi- 
nation of  Ions  —  Displacement  of  One  Ion  :  Electromotive 
Series  of  the  Metals — Destruction  or  Formation  of  a  Com- 
pound Ion  —  Change  in  the  Charges  of  Two  Ions  —  Charge 
and  Discharge  of  Two  Ions,  Electrically — Two  or  More 
Kinds  of  Ionic  Action  Simultaneously. 

NON-IONIC  MODES  OF  FORMING  IONOGENS  :  Acids  and  Bases  — 
Salts. 

XXI.     Sulphur  and  Hydrogen    Sulphide 367 

Sulphur  —  Hydrogen  Sulphide  —  Sulphides  — Poly  sulphides 

—  The  Chemical  Relations  of  the  Element  Sulphur. 

XXII.     The    Oxides   and   Oxygen   Acids  of  Sulphur 378 

Sulphur  Dioxide — Sulphur  Trioxide  —  Oxygen  Acids  of 
Sulphur  —  Sulphuric  Acid  —  Sulphates  —  Constitution  of 
Hydrogen  Sulphate  —  Hyposulphurous  Acid  —  Sulphurous 
Acid  —  Illustration  of  the  Effect  of  Concentration  on  Speed 
of  Interaction — Sulphites  —  Thiosulphuric  Acid  —  Persul- 
phuric  Acid  —  Poly thiouic  Acids  —  Sulphur  Monochloride  — 
Thionyl  Chloride  —  Sulphuryl  Chloride. 

XXIII.  Selenium    and    Tellurium  —  The  Periodic    System 401 

Selenium  — Tellurium  —  The  Chemical  Relations  of  the  Sul- 
phur Family  —  Metallic  and  Non-metallic  Elements  —  Class- 
ification by  Atomic  Weights  —  Mendelejeff 's  Scheme  — 
General  Relations  in  the  System  —  Applications  of  the  Peri- 
odic System. 

XXIV.  Nitrogen  and  its    Compounds  with    Hydrogen  415 

Nitrogen  —  Ammonia  —  Ammonium  Compounds  —  Hydra- 
zine  —  Hydrazoic  Acid  —  Hydroxylamine  —  An  Active  State 
of  Hydrogen  (Nascent  Hydrogen)  —  Halogen  Compounds 
of  Nitrogen. 


XVI 


CONTENTS 


CHAPTEBS  PAGES 

XXV.     The  Atmosphere.     The  Helium  Family  .................       426 

Components  of  the  Atmosphere  —  Air  a  Mixture  —  Liquefac- 
tion of  Gases  —  Liquid  Air  —  Argon  —  Helium  —  Neon, 
Krypton,  and  Xenon. 

XXVI.     Oxides,   and  Oxygen  Acids    of   Nitrogen  ................      438 

Nitric  Acid  —  Nitrates  —  Nitric  Oxide  —  Molecular  Com- 
pounds —  Nitrogen  ^Tetroxide  —  Oxidizing  Actions  of  Nitric 
Acid  —  Nitrous  Acid  —  Nitrites  —  Nitrous  Oxide  —  Explo- 
sives —  The  Principle  of  Transformation  by  Steps. 

XXVII.     Phosphorus  ........................................       455 

Phosphorus  —  The  Electric  Furnace  —  Phosphine  —  Phos- 
phonium  Compounds  —  Halides  of  Phosphorus  —  Oxides  of 
Phosphorus  —  The  Phosphoric  Acids  —  Phosphorous  Acid  — 
Hypophosphorous  Acid  —  Structural  Formulae  of  Salts  of 
Hydrogen  —  Sulphides  of  Phosphorus  —  Comparison  of  Phos- 
phorus with  Nitrogen  and  with  Sulphur. 


XXVIII. 


Carbon  and    the  Oxides  of    Carbon  .....................       473 

Carbon  —  Calcium  Carbide  —  Carbon  Dioxide  and  Carbonic 
Acid  —  Carbonates  —  R61e  of  Chlorophyl-bearing  Plants  in 
Storing  Energy  —  Photochemical  Action  —  Carbon  Monox- 
ide —  Carbonyl  Chloride  and  Urea  —  Carbon  Bisulphide. 

Some  Carbon  Compounds  ...............................       490 

The  Hydrocarbons  —  Petroleum  —  Fractional  Distillation  — 
Methane  —  Organic  Radicals  —  Ethylene  —  Acetylene  — 
Benzene  —  Formic  Acid  —  Acetic  Acid  —  Oxalic  Acid  — 
Carbohydrates  and  Fermentation  —  Alcohols,  Esters,  and 
Ethers  —  Soap  —  Drying  Oils  —  Cyanogen  —  Hydrocyanic 
Acid  —  Cy  anates  —  Thiocyanates. 


XXX.     Flame 


509 


XXXI.     Silicon  and  Boron 518 

Silicon  —  Silicon  Hydride  —  Carbide  of  Silicon  —  Silicon 
Tetrachloride— Silicon  Tetrafluoride  —  Hydrofluosilicic  Acid 
—  Silicon  Dioxide  —  Silicic  Acid  —  Colloidal  Solution  —  Sili- 
cates—  Boron  —  Hydrides  and  Halides  of  Boron  — Boric 
Acid  —  Borates  —  Boron  Trioxide  —  Nitride  and  Carbide. 

XXXII.     The    Base-forming   Elements 530 

Physical  Properties  of  the  Metals  —  General  Chemical  Rela- 
tions of  the  Metallic  Elements  —  Hydrolysis  of  Halogen  Com- 
pounda  Used  to  Distinguish  Metallic  from  Non-metallic 
Elements  —  Salts  of  Complex  Acids  —  Classification  of  the 
Metallic  Elements  by  their  Chemical  Relations  —  Occurrence 
of  th'j  Metallic  Elements  in  Nature  —  Methods  of  Extraction 
from  the  Ores  —  Compounds  of  the  Metals  :  Oxides  and 


CONTENTS 


xvii 


PAGES 


Hydroxides  —  Compounds  of  the  Metals  :  Salts  —  Solubilities 
of  Bases  and  Salts  —  Hydrated  Forms  of  Salts  Commonly 
Used  —  Isomorphism . 

XXXIII.     The  Metals  of  the   Alkalies:    Potassium  and  Ammonium.      548 

The  Chemical  Relations  of  the  Metallic  Elements  of  the  Alka- 
lies —  Potassium  —  the  Hydride  —  Chloride  —  Iodide  —  Bro- 
mide and  Fluorides  —  Hydroxide  —  Oxides  —  Chlorate  — 
Bromate  and  lodate  —  Nitrate  —  Carbonate  —  Cyanide  — 
Sulphate  and  Bisulphate  —  Sulphides  —  Properties  of  Kal- 
ion  :  Analytical  Reactions  —  the  Spectroscope  —  Rubidium 
and  Caesium  —  Ammonium  Chloride  —  Ammonium  Hydrox- 
ide —  Nitrate  —  Carbonate  —  Sulphate  —  Sulphides  —  Micro- 
cosmic  Salt  —  Ammonium  Amalgam  —  Ammonion:  Ana- 
lytical Reactions. 

XXXIV.  Sodium  and  Lithium.    Ionic  Equilibrium  Considered  Quanti- 

tatively         669 

SODIUM  :  Sodium  Hydride  —  Chloride  —  Hydroxide  and  Oxides 
— Nitrate  and  Nitrite  — Carbonate  —  Bicarbonate  —  Sulphate 

—  Thiosulphate  —  Phosphates  —  Tetraborate  —  Silicate  — 
Properties  of  Natrion  :  Analytical  Reactions  —  Lithium. 

IONIC  EQUILIBRIUM  CONSIDERED  QUANTITATIVELY:  The  Simple 
Case — Excess  of  One  Ion  —  Formulation  and  Quantitative 
Treatment  of  the  Case  of  Excess  of  One  Ion  —  Special  Case 
of  Saturated  Solutions  —  Illustration  of  the  Principle  of  Ion- 
Product  Constancy  —  Other  Illustrations — Exercises. 

XXXV.  The  Metals  of  the  Alkaline  Earths 688 

The  Chemical  Relations  of  the  Elements  —  Calcium  and  its 
Compounds— The  Phase  Rule,  a  Method  of  Classifying  All 
Systems  in  Equilibrium  —  Hard  Water—  Mortar  and  Cement 

—  Theory  of  Precipitation —  Rule  for  Solution  of  Substances 

—  Interaction  of  Insoluble   Salts  with  Acids,  Resulting  in 
Solution  of  the  Salt  —  Glass  —  Calcion :  Analytical  Reactions 

—  Strontium  and  its  Compounds  —  Barium  and   its  Com- 
pounds —  Analytical  Reactions  of  the  Calcium  Family. 

XXXVI.    Copper,   Silver,  Gold 614 

The  Chemical  Relations  of  the  Copper  Family  — 
Copper  and  its  Compounds  —  The  Solution  of  In- 
soluble Salts  when  Complex  Ions  are  Formed  —  Silver  and 
its  Compounds  —  Electro-plating  —  Photography  —  Gold  and 
its  Compounds. 

XXXVII.      Glucinum,    Magnesium,    Zinc,    Cadmium,    Mercury.      The 

Recognition  of   Cations  in  Qualitative  Analysis 641 


XV111 


CONTENTS 


CHAPTKB8 

XXXVIII. 


Electromotive  Chemistry  ...............................      664 

Factors  and  Units  of  Electrical  Energy  —  Displacement  Cells 

—  Potential    Differences    Produced   by  the    Metals  Singly 

—  Application    to    Cells  —  Other   Applications  :     Couples  : 
Concentration   Cells  —  Electrolysis  :   Discharging  Potentials 

—  The  Factors  of  Energy  —  Methods  of  Measuring  Chemical 
Activity. 


XXXIX.     Aluminium  and  the  Metals  of  the  Earths 

The  Rare  Elements  of  This  Family  —  Aluminium  and  its 
Compounds  —  Dyeing  :  Mordanting  —  Kaolin  and  Clay  : 
Earthenware  and  Porcelain. 


XL.    Germanium,  Tin,  Lead 


XLI.     Arsenic,  Antimony,  Bismuth 


681 


692 
707 


XLII.     The  Chromium  Family.    Radium 722 

The  Chemical  Relations  of  the  Family  —  Chromium  — 
Derivatives  of  Chromic  Acid  —  Chromic  Compounds  — 
Chromous  Compounds  —  Analytical  Reactions  —  Molyb- 
denum —  Tungsten  —  Uranium  —  Radium  —  The  Discovery 
of  the  Element  —  Properties  of  Radium  Compounds  — 
The  Radiations  of  Radium  Salts  —  The  Decay  of  an  Element. 

XLIII.     Manganese 737 

XLIV.    Iron,  Cobalt,  Nickel 746 

763 


XLV.     The  Platinum  Metals 

Ruthenium  and   Osmium  —  Rhodium  and  Iridium 
dium  and  Platinum. 


Palla- 


INTRODUCTION   TO 
GENERAL    INORGANIC    CHEMISTRY 


INTRODUCTION    TO 
GENERAL   INORGANIC   CHEMISTRY 


CHAPTER   I 
INTRODUCTORY  I 

HUMAN  knowledge  has  become,  in  recent  times,  so  extensive  and 
complex  that  the  truths  ascertained  have  had  to  be  divided,  more  or 
less  arbitrarily,  into  groups.  Thus,  the  study  of  animals,  their  classi- 
fication by  structure,  life-history,  distribution,  and  so  forth,  form  the 
group  known  as  zoology.  Such  a  group  is  called  a  science,  and  in- 
cludes a  more  or  less  distinct  body  of  knowledge.  That  the  boun- 
daries of  these  groups  are  purely  arbitrary,  and  do  not  exist  in  the 
subject-matter  itself,  is  seen  at  once  in  our  own  treatment  of  them. 
Thus,  for  convenience,  we  take  the  structure  of  animals  by  itself  and 
style  it  anatomy;  but  we  include  in  the  science  of  physiology  the 
study  of  the  way  in  which  the  parts  of  both  plants  and  animals  per- 
form their  functions ;  and  we  assemble  cognate  parts  of  two  groups 
in  sciences  like  astro-physics  and  physical  chemistry.  The  sciences, 
therefore,  are  not  mutually  exclusive,  and  their  boundaries  overlap  in 
every  direction. 

The  difficulty  in  deciding  what  are  the  most  convenient  boundaries 
is  as  great  with  chemistry  as  with  the  other  groups.  At  the  one 
extreme  we  have  the  abstract  sciences,  logic  and  mathematics.  At 
the  other  extreme  lie  the  concrete  sciences  like  geology,  zoology,  and 
astronomy.  The  former  are  not  concerned  primarily  with  the  study 
of  matter  at  all,  but  with  that  of  abstract  conceptions.  The  latter 
deal  with  definite  aggregates  of  matter,  such  as  the  nature  and  his- 
tory of  a  particular  deposit  of  sulphur  and  their  relation  to  those 
of  other  deposits  of  sulphur,  or  the  structure  and  history  of  a  partic- 
ular collection  of  organic  material  known  as  a  pike,  and  their  relation 
to  the  structure  and  history  of  a  mass  of  similar  material  called  a 
salmon.  Between  these  two  sets  of  sciences  are  the  regions  occupied 
by  the  abstract-concrete  sciences,  physics  and  chemistry.  These 
sciences  deal  in  part  with  the  same  portions  of  matter,  but  in  a  more 


4  INORGANIC   CHEMISTRY 

abstract  way  than  do  geology  or  biology.  To  them,  all  specimens  of 
pure  sulphur  are  alike,  whether  they  have  been  formed  by  volcanic 
action,  or  have  been  deposited  by  bacteria  in  an  entirely  different 
manner.  In  particular,  the  line  which  divides  physics  and  chemistry 
from  one  another  is  often  difficult  to  draw.  It  is  assumed,  however, 
that  the  reader  is  already  familiar  with  the  elements  of  physics,  and 
so,  in  place  of  entering  upon  an  academic  discussion  of  the  nature  of 
this  line,  we  shall  allow  its  location,  to  emerge  as  we  proceed. 

The  same  principle  of  grouping  is  pursued  within  each  field. 
Thus  the  preparation  and  properties  of  chemical  compounds  is  called 
descriptive  chemistry,  and  the  content  of  this  portion  of  the  subject 
is  in  turn  classified  according  to  a  plan  involving  the  consideration  of 
the  constituents  of  each  compound.  The  study  of  the  proportions  of 
the  constituents  in  compounds,  of  the  conditions  under  which  chem- 
ical action  occurs,  and  related  matters  of  a  more  abstract  character, 
are  grouped  together  in  theoretical  chemistry,  which  is  likewise  sub- 
divided. Again,  the  means  that  have  been  devised  for  recognizing 
the  components  of  mixtures  or  compounds  and  measuring  their  quan- 
tities constitute  the  several  branches  of  analysis.  The  subdivisions 
of  chemistry  of  this  kind  are  numerous. 

The  ideal  in  view  in  thus  classifying  the  content  of  a  science  is  to 
convert  it  into  an  organized  body  of  knowledge.  The  various  ways 
used  to  organize  the  facts  of  a  science  will  be  presented  in  detail  as 
opportunity  offers.  These  ways  constitute  what  is  called  the  scientific 
method. 

It  is  only  by  following  intelligently  the  way  in  which  the  science 
is  manufactured,  step  by  step,  out  of  the  raw  material  furnished  by 
observation  and  experiment,  that  the  student  can  gain  a  sound  foun- 
dation for  more  advanced  work  in  the  same  science,  or  a  mental  train- 
ing broad  enough  in  -its  tendency  to  add  measurably  to  his  efficiency 
in  every  other  task.  The  chief  object  of  useful  thought,  no  matter 
whether  the  problem  is  one  of  language,  history,  business,  or  life,  is 
to  organize  isolated  facts  into  knowledge,  and  the  means  of  success- 
fully accomplishing  this  is  the  use  of  the  scientific  method. 

Chemical  Phenomena:  Their  Most  Obvious  Character- 
istic. —  Chemistry  being  primarily  an  experimental,  and  not  a  purely 
abstract  science,  we  can  best  make  our  first  approach  to  it  through  the 
consideration  of  some  familiar  chemical  phenomena  which  shall  serve 
as  illustrations. 


INTRODUCTORY   I  5 

When  clean  iron  is  exposed  to  air  and  moisture  it  becomes  rusty, 
first  on  the  surface,  and  finally  throughout  its  whole  mass.  The  iron 
with  which  we  start  is  a  dark-gray,  metallic-looking  substance,  which 
has  a  high  specific  gravity  (about  7.5),  is  ductile,  has  great  tenacity, 
and  is  readily  attracted  by  a  magnet.  Kust,  on  the  other  hand,  is  a 
reddish  or  brownish  substance  which  has  an  earthy  appearance,  is 
much  lighter  specifically  than  iron  (sp.gr.  about  4.5),  is  brittle,  and 
is  not  attracted  by  a  magnet.  The  phenomenon  seems  to  have  con- 
sisted in  the  gradual  substitution  of  the  second  of  these  substances 
for  the  first. 

The  chemical  fact  here  is  that  iron,  when  in  contact  with  air  and 
moisture,  gives  rust.  Consideration  will  show,  however,  that  we  do 
not  perceive  this  fact  except  by  noting  the  physical  qualities  of  the 
original  and  of  the  final  materials.  The  data  describing  a  chemical 
phenomenon  consist  in  an  enumeration  of  physical  observations,  like 
the  above  physical  properties  of  iron  and  rust.  It  is,  indeed,  an  in- 
variable characteristic  of  every  chemical  phenomenon  that  our  informa-- 
tion  is  all  derived  from  physical  study  of  the  materials.  Thus,  when 
a  candle  burns,  it  undergoes  a  chemical  change,  and  we  observe  a 
solid,  waxy  substance  disappearing  progressively  from  view.  A  closer 
study  of  the  facts  shows  that  a  mixture  of  gases  is  rising  from  the 
flame,  and  we  find  that  the  material  of  the  candle  is  contained  in  this. 
Here,  again,  we  use  physical  means  of  observation. 

It  is  further  true  of  these,  as  of  all  chemical  phenomena,  that  the 
physical  properties  of  the  original  and  those  of  the  final  substances 
are  invariably  entirely  different.  We  observe  also  that  this  sort  of 
transformation  is  always  abrupt.  As  the  change  spreads,  the  minute 
fragment  of  iron  of  one  moment  becomes  the  minute  fragment  of  rust 
of  the  next,  and  is  endowed  at  once  with  all  the  new  properties  in  full 
measure.  No  part  of  the  mass  loses  its  magnetic  qualities  completely, 
for  example,  before  it  has  reached  the  limit  of  change  in  color  or  ten- 
acity. Since  we  are  compelled  to  define  the  species  of  matter  by 
their  constant  physical  properties  (see  p.  35),  the  possession  by  two 
or  more  bodies  of  permanently  different  properties  constitutes  them 
ipso  facto  samples  of  different  materials.  We  may,  therefore,  say, 
more  briefly,  that  the  products  of  a  chemical  change  are  recognized 
at  once  to  be  new  and  different  substances. 

Thus  the  moat  obvious  characteristic  of  a  chemical  phenomenon  is 
that  all  the  physical  properties  of  the  substances  alter,  that  this  alter- 
ation is  abrupt,  that,  in  fact,  the  products  are  different  substances,  and 


6  INORGANIC  CHEMISTRY 

that  the  recognition  and  study  of  such  a  phenomenon  is  accomplished 
entirely  by  observations  of  a  physical  nature.  Other  characteristics, 
less  obvious,  but  equally  constant,  will  presently  be  brought  to  light. 

The  chemical  phenomena  which  are  familiar  are  innumerable. 
The  burning  of  illuminating-gas  or  of  kerosene,  the  coagulation  of 
the  white  of  an  egg,  the  decay  of  animal  and  vegetable  matter,  the 
action  of  hard  water  on  soap,  the  "burning"  of  limestone,  and  the 
slaking  of  quicklime,  are  examples.  The  reader  should  analyze  these 
changes,  applying  the  above  criteria,  and  see  for  himself  why  they 
are  classed  as  chemical. 

So  far  as  this  first  characteristic  is  concerned,  some  merely  physi- 
cal phenomena  will  be  recalled,  which  are  not  unlike  the  rusting  of 
iron.  Thus,  when  water  is  raised  in  temperature,  its  properties  begin 
to  alter :  it  becomes  more  mobile,  its  specific  gravity  changes,  its 
refracting  power  for  light  is  modified,  and  so  forth.  At  first,  how- 
ever, these  changes  proceed  by  imperceptible  gradations  and  lack  the 
abruptness  of  chemical  change.  But  when  100°  C.  is  reached  the 
water  passes  into  steam,  and  the  whole  of  the  properties  of  this  gas 
are  different  from  those  of  water.  Conversely,  by  cooling,  the  water 
may  be  converted  into  ice,  and  again  there  is  an  abrupt  and  profound 
change  in  properties.  On  this  account,  some  chemists  would  classify 
"changes  of  state"  as  chemical  phenomena.  More  usually,  however, 
the  absence  of  the  other  characteristics  yet  to  be  mentioned  is  held  to 
justify  the  more  common  view,  and  ice,  water,  and  steam  are  regarded 
as  identical,  and  not  different  substances. 

The  very  language  we  use  bears  testimony  to  the  universal  ac- 
ceptance of  the  view  that  a  physical  change  does  not  constitute  a 
fundamental  alteration.  Thus,  we  have  solid  lead  and  molten  lead, 
air  (the  gas)  and  liquid  air.  In  the  case  of  water  alone  has  it  been 
found  convenient  to  distinguish  ice,  water,  and  steam  by  separate  names. 

Fact  and  Generalization  or  Laiv.  — If  the  preceding  paragraph 
be  reexamined,  it  will  be  seen  that  a  number  of  facts  are  mentioned 
in  it.  We  begin  the  organization  of  knowledge  according  to  the 
scientific  method  by  trying  to  determine  the  facts.  Thus,  we  find 
some  specimens  of  iron  are  variously  colored,  and  some  are  brittle. 
Examination  shows,  however,  that  the  former  peculiarities  are  due  to 
paint,  for  example,  and  the  latter  to  the  presence  of  carbon  and  other 
foreign  materials  in  the  iron  (cast  iron).  Finally,  we  ascertain  the 
facts  that  iron  itself  is  gray  and  tough. 


INTRODUCTORY  I  7 

Facts  are  the  ultimate  units  of  the  structure  of  a  science.  Thus,  a  single  iso- 
lated observation,  no  matter  how  accurately  made,  does  not  furnish  us  with  the 
sort  of  fact  that  can  receive  a  permanent  place  in  our  collection.  It  is  only  after 
much  research  and  thought  that  we  can  ascertain  which  are  the  fixed  elements  in 
the  variety  of  experience  in  any  line,  and  so  determine  what  are  the  facts,  in  the 
sense  in  which  we  have  used  the  term. 

Putting  together  a  statement  like  that  appearing  in  heavy  type 
above,  is  the  second  step.  We  examine  facts  of  a  like  kind,  or  per- 
taining to  like  phenomena,  to  see  whether  any  general  statement  can 
be  made  that  will  cover  some  feature  common  to  the  whole  of  them. 
In  the  above  illustrations,  after  settling  the  intrinsic  properties  of 
several  substances,  and  then  determining  the  facts  about  the  way  in 
which  these  properties  are  affected  under  certain  circumstances,  we 
decide  that,  when  all  the  properties  change  abruptly  in  a  permanent 
way,  the  cases  in  which  this  takes  place  shall  constitute  a  distinct 
class,  and  we  call  them  cases  of  chemical  change.  The  statement 
which  in  a  few  words  or  phrases  sums  up  those  features  of  all  the 
phenomena  of  like  kind  which  are  constant,  is  called  a  generalization. 
Often  it  is  called  a  law :  there  is  no  decided  difference  in  the  usage 
of  the  two  words.  Sometimes  it  is  described  as  a  principle  of  the 
science.  A  generalization  or  law  in  chemistry,  therefore,  is  a  brief 
statement  describing  some  constant  mode  of  behavior. 

Of  course,  the  same  set  of  facts  may  be  viewed  in  many  ways.  Picking  out 
the  relationships  which  are  most  comprehensive  and,  at  the  same  time,  are  best 
fitted  to  form  part  of  still  broader  generalizations,  or  to  take  their  places  alongside 
of  equally  broad  ones,  requires  the  highest  ability.  The  most  important  laws, 
like  that  describing  the  behavior  of  gases  when  compressed  (Boyle's  law),  are 
usually  connected  with  the  names  of  the  men  by  whom  the  relationships  were  dis- 
covered and  the  generalizations  formulated. 

The  word  "formulate,"  as  applied  to  a  law,  is  preferable  to  the  word  "dis- 
cover." The  latter  is  ambiguous  and  suggests  that  the  law  existed  before  it  was 
found.  Even  the  relation  which  it  puts  into  words,  did  not,  properly  speaking, 
exist,  because  relations  are  picked  out  of  a  complex  by  the  mind,  and  the  particu- 
lar relation  selected  is  a  property  of  the  mind  and  not  of  the  constituents  of  the 
complex  itself. 

The  reader  must  beware  of  the  misconception  that  a  law  enforces  behavior  in 
accordance  with  its  tenor.  The  mere  statement  that  every  piece  of  matter  attracts 
every  other,  cannot  compel  a  stone  to  fall ;  nor  can  the  mode  of  behavior  of  one 
falling  stone  persuade  any  other  to  do  likewise.  The  law  is  simply  a  record 
of  what  is  invariably  observed  to  happen. 

It  is,  therefore,  also  very  misleading  if  we  permit  ourselves  to  say  that  Boyle's 
law  "  acts"  so  as  to  "cause"  gases  to  behave  in  a  certain  way,  or  that  the  law 
"operates"  to  "produce"  a  certain  behavior,  or  that  other  behavior  is  "impos- 
sible" to  the  gas,  or  that  the  law  of  cohesion  "  intervenes  "  when  the  gas  is  under 


8  INORGANIC   CHEMISTRY 

low  pressure,  and  causes  its  behavior  to  "diverge  from  that  required  "  by  Boyle's 
law,  or  to  say  that  a  gas  "  disobeys"  Boyle's  law.  In  scientific  discussions,  such 
figures  of  speech  are  alone  permissible  as  throw  light  on  the  subject.  Phrases 
like  the  above,  common  as  their  use  is,  have  been  selected  apparently  with  a  view 
to  introducing  a  maximum  of  distortion  and  obscurity.  It  is  the  gas  that  "  acts  " 
and  gives  rise  to  the  making  of  Boyle's  law,  and  the  latter  is  only  an  epitome  of 
the  way  it  acts.  Everything  that  may  be  conceived  may  also  be  possible  in  nature, 
although  there  are  many  things  which  have  not  yet  been  known  to  occur.  There- 
fore, we  may  not  say  that  it  is  "  impossible  "  for  a  gas  to  alter  its  volume  otherwise 
than  inversely  with  the  pressure.  The  progress  of  science  would  be  almost  com- 
pletely arrested,  if,  every  time  we  succeeded  in  formulating  a  seemingly  satisfac- 
tory statement  of  truth  in  regard  to  some  set  of  phenomena,  the  exhibition  by 
nature  of  any  behavior  which  was  in  conflict  with  our  statement  became  forthwith 
"impossible."  It  is  not  the  gas  which  "diverges"  from  our  statement  or  "dis- 
obeys" our  law,  but  our  statement  which  is  proved  by  the  behavior  of  the  gas  to 
be  inaccurate.  Our  procedure,  in  such  cases,  is  always  more  logical  than  our  lan- 
guage, for  we  never  attempt  to  cure  the  gas  of  its  error,  but  always  the  law  itself 
by  suitable  modification  in  its  phraseology. 

The  Explanation  of  Rusting.  —  In  considering  the  complete 
disguise  which,  is  assumed  by  a  substance  that  has  undergone  chemi- 
cal change,  the  first  question  which  arises  is :  Can  we  in  any  way 
account  for  this  great  and  permanent  change  in  properties  which  very 
generally  distinguishes  the  chemical  phenomenon  from  the  physical  ? 
Some  additional  facts  will  be  required  before  we  can  answer  this 
question.  Many  attempts  have  been  made,  from  the  earliest  times,  to 
learn  the  exact  nature  of  the  change  known  as  rusting.  It  was  found 
that  many  metals  besides  iron  underwent  a  somewhat  similar  altera- 
tion, and  that,  where  the  transformation  did  not  occur  spontaneously, 
heating  prompted  it.  The  first  fact  which  seemed  to  throw  light  on 
the  subject  was  the  observation  that  a  piece  of  metal  becomes  heavier 
when  it  rusts.  This  was  noticed  as  early  as  1630,  by  a  French  physi- 
cian, Jean  Key.  His  work  was  done  chiefly  with  lead  and  tin,  the 
former  of  which  gives  a  dirty  yellow,  and  the  latter,  a  white  powder, 
on  rusting.  He  inferred  correctly  from  his  experiments  that  contact 
with  the  air  had  something  to  do  with  this  chemical  change.  Other 
investigations  on  the  same  subject  were  made  by  Boyle  (1627-1691), 
Mayow  (1645-1679),  and  Hooke  (1635-1703),  in  England,  and  they  led 
to  the  same  conclusion.  The  increase  in  weight  was  to  be  accounted 
for,  therefore,  by  the  supposition  that  some  material  from  the  air  had 
been  added  to  the  metal  during  the  process.  In  other  words,  iron,  for 
example,  was  one  substance  composed  of  iron  only,  and  rust  was 
another  substance  composed  of  iron  and  some  other  material  taker 


INTRODUCTORY  I  9 

from  the  atmosphere;  and,  the  two  substances  being  different  in 
composition,  their  properties  might  naturally  be  expected  to  differ. 
The  substance  taken  from  the  air  was  subsequently  named  oxygen. 

A  rough  imitation  of  the  rusting  of  iron  may  be  shown  to  be  accompanied  by 
an  increase  in  weight.  Iron  powder  is  suspended  by  means  of  a  magnet  over  one 
pan  of  a  balance,  and  the  equipoise  is  restored  by  placing  small  shot  on  the  other 
pan.  When  the  iron  is  heated,  union  with  oxygen  begins,  and,  after  a  time,  the 
pointer  inclines  markedly  to  one  side.  The  product  here  is  magnetic  oxide  of 
iron  (Fe304),  however,  and  not  rust  (Fe203,  icH2O). 

An  Older  View :  Phlogiston.  —  Simple,  and  to  a  certain  extent 
satisfactory,  as  this  explanation  appears  to  us,  it  must  be  said  that  it 
gained  little  sympathy  from  the  contemporaries  of  these  men.  The 
other  investigators  had  been  prejudiced  by  a  remarkable  hypothesis 
which  was  supposed  to  explain  both  rusting  and  combustion.  Starting 
with  a  suggestion  of  Plato's,  that  during  combustion  some  material 
escaped  from  the  burning  body,  and  that  the  flames  and  heat  repre- 
sented the  vigor  with  which  this  substance  rushed  out,  Stahl  and 
Becher  invented  the  idea  that  a  substance,  which  they  called  "  phlo- 
giston," was  contained  in  all  materials  capable  of  rusting  or  burning. 
Its  escape  accounted  for  the  phenomena  of  combustion,  and  its  absence 
for  the  alteration  in  properties  of  the  residual  substance.  They  were 
perfectly  aware  that  the  material  was  heavier  after  rusting  than 
before,  but  refused  to  sacrifice  their  hypothesis  to  a  mere  fact  like 
this.  So  they  ingeniously  appended  the  suggestion  that  phlogiston 
was  a  substance  which  not  only  was  not  subject  to  gravitation,  but 
possessed  the  opposite  property  of  levity.  Thus,  its  escape  rendered 
the  material  from  which  it  issued  heavier  than  before.  Instead  of 
demanding  the  preparation  and  examination  of  phlogiston  itself,  and 
the  demonstration  that  it  weighed  less  than  nothing,  the  generality  of 
chemists  of  that  age  accepted  the  idea  without  proof.  It  is  not  sur- 
prising, therefore,  that  many  of  their  attempts  to  explain  chemical 
phenomena  on  the  basis  of  an  arbitrary  assumption  like  this  should 
have  proved  confusing  and  infertile.  The  fact  that  foreign  matter 
was  actually  gained  by  a  body  during  the  process  of  rusting  was  not 
generally  accepted  until  it  was  demonstrated  anew  by  Lavoisier 
(1774).  He  showed  that  a  portion  of  the  air  really  disappeared  when 
tin  rusted,  and  that  the  increase  in  weight  of  the  tin  corresponded 
with  the  loss  in  matter  of  the  air.  The  whole  development  of  chem- 
istry was  stunted  by  the  general  belief  in  the  conception  of  phlogiston 


10  INORGANIC   CHEMISTRY 

and,  during  the  one  hundred  and  fifty  years  which  passed  between 
Jean  Rey's  discovery  and  Lavoisier's,  relatively  little  progress  was 
made. 

The  introduction  of  the  balance  into  the  chemical  laboratory,  and  the  first 
use  of  measurements  of  weight  as  a  means  of  exploring  and  explaining  chemical 
changes,  is  frequently  ascribed  to  Lavoisier.  As  a  matter  of  fact,  it  is  difficult  to 
state  when  measurement  of  weight  first  became  the  chief  ally  of  the  chemist  in  his 
work.  Important  and  conclusive  results  were  obtained  by  its  means,  however, 
before  the  time  of  Lavoisier,  for  example,  by  Jean  Rey,  Boyle,  and  Black  (1728- 
1799). 

Explanation  in  Science.  —  The  word  explanation,  which  was 
employed  repeatedly  in  the  last  two  sections,  is  used  in  science  as  the 
name  of  a  definite  process.  It  stands  simply  for  a  description  in 
greater  detail.  Thus,  when,  to  the  acquaintance  with  the  outward 
manifestations  of  rusting,  we  are  able  to  add  the  further  description 
that  it  is  produced  by  the  union  of  oxygen  from  the  air  with  iron,  we 
feel  increased  satisfaction,  and  we  say  that  an  explanation  has  been 
found.  Sometimes  we  get  this  satisfaction  by  an  explanation  which 
is  a  description,  not  of  additional  facts,  but  of  some  imaginary  con- 
dition which  takes  their  place  in  our  thought.  The  hypothesis  (or 
supposition)  of  phlogiston  was  somewhat  of  this  nature  (see  Formula- 
tive  hypothesis). 

There  is  no  such  thing  as  a  final  explanation.  At  the  very  next 
step,  when  we  ask  why  the  union  of  oxygen  and  iron  produces  a  body 
that  is  red  and  non-magnetic,  we  are  compelled  to  say  we  do  not  know. 
Even  a  supposition  in  regard  to  how  this  happens  is  as  yet  lacking. 

An  explanation  in  science  never  professes  for  a  moment  to  give  the 
reasons  for  any  occurrence.  We  simply  do  not  know  why  behavior  in 
nature  is  as  it  is.  Hence  statements  like  that  in  regard  to  union  with 
oxygen,  constitute  all  the  answer  that  we  can  give  to  the  question  with 
which  a  recent  section  (p.  8)  opened. 

Three  Illustrative  Chemical  Phenomena.  —  Since  oxygen  is 
an  invisible  gas,  the  demonstration  that  rusting  consists  in  the  union 
of  this  gas  with  a  metal,  requires  somewhat  complicated  apparatus. 
The  next  illustration,  while  lacking  historical  interest,  is  simpler,  be- 
cause both  substances  are  visible,  and  are  easily  handled. 

Iron  unites,  not  only  with  oxygen  from  the  air,  but  also,  with  almost 
equal  ease,  with  sulphur.  To  study  the  behavior  of  the  materials  we 
must  know  their  properties.  The  properties  of  iron  we  have  already 


INTRODUCTORY  I 


11 


FIG.  1. 


enumerated  (p.  5).  Sulphur  is  a  pale-yellow  substance  of  low  spe- 
cific gravity  (sp.  gr.  2).  It  is  easily  melted  (m.-p.  115°  C).  and,  al- 
though it  does  not  dissolve  in  water,  it  may  be  dissolved  completely  in 
carbon  disulphide  (41 : 100  at  18°).*  It  crystallizes  in  rhombic  forms 
(Fig.  1).  Now,  when  iron  filings  and  powdered  sulphur  are  rubbed 
together  in  a  mortar,  the  product,  although  it  has  a  different  color  from 
either  of  the  constituents,  is  still  really  com- 
posed of  the  two  original  kinds  of  matter,  side 
by  side.  With  the  help  of  a  microscope  and 
a  needle,  they  can  be  picked  apart  completely. 
By  manipulation  of  the  mixture  with  a  mag- 
net, we  may  remove  some  of  the  iron  without 
much  difficulty. 

Again,  using  the  solubility  of  sulphur  in 
carbon  disulphide  and  the  insolubility  of  iron, 
we  may  shake  the  mixture  with  this  liquid  in 

a  test-tube,  and  so  dissolve  out  the  sulphur  (Fig.  2).  When  the  con- 
tents of  the  tube  are  poured  on  to  a  filter  (Fig.  3),  the  liquid  (the 
filtrate)  runs  through,  carrying  all  dissolved  matter  with  it,  and  leav- 
ing undissolved  matter  behind.  The 
former  may  be  allowed  to  evaporate 
(Fig.  4),  when  yellow  crystals  of  rhom- 
bic outline  will  be  found  to  be  the  sole 
residue.  The  dark  material  remaining 
on  the  filter,  when  dry,  is  wholly  at- 
tracted by  a  magnet.  All  these  facts 
convince  us  that  the  properties  of  the 
components  are  still  unaltered,  and  that, 
therefore,  no  chemical  change  has  oc- 
curred. 

If  we  now  put  some  of  the  original 
mixture  into  a  test-tube  and  warm  it, 
we  soon  notice  a  rather  violent  develop- 
ment of  heat  taking  place,  the  contents 

begin  to  glow,  and  what  appears  to  be  a  form  of  combustion  spreads 
through  the  mass.  The  heating  employed  at  the  start  falls  far 
short  of  accounting  for  the  much  greater  heat  produced.  When 
these  phenomena  have  ceased,  and  the  test-tube  has  been  allowed  to 

*  This  expresses  the  fact  that  41  parts,  by  weight,  of  sulphur  dissolve  in  100 
parts,  by  weight,  of  carbon  disulphide  at  18°  C. 


FIG.  2. 


12 


INORGANIC   CHEMISTRY 


cool,  we  find  that  it  now  contains  a  somewhat  porous-looking,  black 
solid.  This  material  is  brittle ;  it  is  not  magnetic ;  it  does  not  dissolve 
in  carbon  disulphide ;  and  close  examination,  even  under  a  microscope, 
does  not  reveal  the  presence  of  different  kinds  of  matter.  This  sub- 
stance is  known  to  chemists 
as  ferrous  sulphide,  and,  as 
we  see,  its  properties  are 
entirely  different  from  those 
of  the  constituents. 

A  substance  formed  in 
this  way  by  the  union  of 
other  materials  is  called  a 
compound.  The  rusts  given 
by  various  metals  are  there- 
fore compounds  also. 

The  second  illustration 
is  selected  on  account  of  its 
historical  interest.  One  of 
the  earliest  chemical  changes 
in  which  a  gas,  recognized 
to  be  distinct  from  air,  was 

observed  among  the  products,  was  noticed  by  Priestley  (1774).*  The 
observation  was  made  with  mercuric  oxide,  a  bright  red,  rather 
heavy  powder.  When  this  substance  is  heated  (Fig.  5),  we  find  that 
a  gas  is  given  off,  which  is  easily  shown  to  be  different  from  air,  since 
a  glowing  splinter  of  wood  is  instantly  relighted  on  being  immersed 
in  it.  The  gas  is  pure  oxygen.f 
We  notice  also  during  the  heating 
that  a  sort  of  mirror  appears  on 
the  sides  of  the  tube.  As  this 
shining  substance  accumulates  it 
takes  the  form  of  globules,  which 
may  be  scraped  together.  It  is, 

in  fact,  the  metal  mercury,  or  quicksilver.  If  the  heating  continues 
long  enough,  the  whole  of  the  powder  eventually  disappears,  and  is 
converted  into  these  products. 

*  An  English  nonconformist  minister  who  occupied  his  leisure  time  with  exper- 
iments in  chemistry.  He  afterwards  moved  to  the  United  States,  and  died  m 
Northumberland,  Pa. 

t  Air  (g.TJ.)  is  a  mixture  of  which  only  one-fifth  is  oxygen. 


FIG.  3. 


FIG.  4. 


INTRODUCTORY   I 


The  extensive  nature  of  the  change  in  properties  in  this  case  is  evi- 
dent.    It  should  also  be  observed  that  continuous  heating  is  required 
to  maintain  this  change  in  operation.    It  differs  markedly  from  the  iron 
and  sulphur  case  in  this  respect.    When  the  flame  is  removed,  the  evo- 
lution of  oxygen  ceases.     The  significance  of  this  will  appear  shortly. 
The  third  and  last  example  is  taken  purposely  in  order  to  illustrate 
the  variety  of  ways  in  which  chem- 
ical change  may  be  carried  out.     It 
is  the   interaction  of  silver  nitrate 
and  sodium  chloride  (common  salt). 
The  substances  may  be  recognized 
by  the  form  of  the  crystals  of  which 
they  consist.    The  latter  is  composed 
of  small  cubes  (Fig.   6),  while  the 
former  presents  a  less  familiar  form 
geometrically  (Fig.  7).      Both  sub- 
stances  are    capable   of    being    dis- 
solved in  water  and,  for  this  experi-  FIG.  5. 
ment,  portions  of  each  substance  are 

shaken  in  separate  vessels  with  water,  until  none  of  the  solid  remains. 
When  the  solutions  are  now  poured  together,  we  observe  that  the 
clear  liquids  at  once  become  opaque,  and  that  a  dense 
mass  of  white,  solid  material  appears  suspended  in 
the  mixture  (Fig.  8).     This  white  substance  consists 
of  an  extremely  fine  powder  without  any  observable 
crystalline   form.     We  know   at  once  that   it  must 
represent  a  new  substance,  since  it  would  not  have 
FIG.  6.  appeared  had  it  been  soluble  in  water  like  the  two 

materials  from  which  it  was 
made.  We  continue  adding  the  one  liquid  gradu- 
ally to  the  other  until  no  further  formation  of 
this  solid  takes  place,  and  then  stop.  By  filtra- 
tion (Fig.  3),  we  obtain  the  insoluble  material 
(the  precipitate)  upon  the  filter  paper,  and  the 
clear  liquid  (the  filtrate)  passes  through  and  is  caught  in  the  vessel 
below. 

We  are  confronted  with  two  possibilities  :  either  both  the  original 
materials  have  come  together  to  form  one  white  insoluble  material,  or 
some  other  product  (or  products)  may  be  present  in  addition  to  it.  In 
the  latter  case,  search  must  evidently  be  made  in  the  liquid.  By  evap- 


FIG.  7. 


14 


INORGANIC  CHEMISTRY 


FIG.  8. 


orating  the  filtrate  in  a  suitable  vessel  (Fig.  4),  we  find  that  the  second 
assumption  represents  the  fact,  for  a  considerable  quantity  of  a  white 
crystalline  substance  remains.  The  homogeneous  character  of  this 
shows  that  there  was  but  one  product  in  solution,  while  the  same  prop- 
erty of  the  precipitate  shows  that  there  are  but  two  products  altogether. 
The  insoluble  material  is  composed  of  silver 
and  chlorine,  and  it  is  known  as  silver  chloride. 
Like  some  other  compounds  of  silver,  it  darkens 
on  exposure  to  light,  turning  first  purple  and  then 
brown,  and  being  decomposed  by  this  agency  into 
its  constituents.  The  soluble  solid  obtained  from 
the  filtrate,  we  recognize  as  identical  with  a 
mineral,  sodium  nitrate,  which  is  found  in  Peru. 
Its  crystals  are  rhombohedral  (Fig.  9).  They 
resemble  cubes  which  have  been  slightly  dis- 
torted by  pressing  inwards  two  opposite  corners. 
A  strict  investigation  of  all  four  substances 
shows  that  the  following  statement  presents  the 
facts  in  regard  to  the  nature  of  the  change :  —  Nitrate  of  silver, 
consisting  of  silver,  nitrogen,  and  oxygen,  with  sodium  chloride,  con- 
sisting of  sodium  and  chlorine,  have  changed 
into  silver  chloride,  consisting  of  silver  and  chlo- 
rine, and  sodium  nitrate,  consisting  of  sodium, 
nitrogen,  and  oxygen.  This  change  presents 
several  features  which  distinguish  it  from  the 
previous  ones :  it  is  much  more  complex ;  it  takes 
place  in  the  presence  of  water ;  it  requires  no 
heating  for  its  promotion  ;  and  the  change  is  complete  the  instant  the 
materials  have  been  mixed,  while  the  others  required  a  good  deal  of 
time  for  their  accomplishment. 


The  Kinds  of  Chemical  Change.  —  Having  these  three  cases 
before  us,  —  and  they  are  types  of  most  chemical  phenomena,  —  we  now 
proceed  to  analyze  them,  and  so  take  another  step  towards  explaining 
(i.e.,  describing  in  detail)  the  nature  of  a  chemical  phenomenon.  In 
the  iron  and  sulphur  experiment  two  materials  were  used,  and  a  dif- 
ferent one  with  new  properties  was  produced.  Here,  in  chemical 
language,  the  first  two  substances  united  or  underwent  combination. 
The  rusting  of  iron,  lead,  and  tin  belongs  also  to  this  class.  In  the 
second  illustration  one  material  was  used,  and  it  was  driven  apart  by 


FIG.  9. 


INTRODUCTORY  I  15 

heating  so  that  two  new  ones  arose.  A  chemical  phenomenon  of  this 
kind  is  called  a  decomposition.  The  last  was  the  most  complex,  but 
a  little  examination  shows  that  it  does  not  present  any  novel  features. 
The  statement  we  gave  (p.  14)  about  the  way  in  which  the  constit- 
uents were  distributed,  before  and  after  the  change,  shows  that  the 
original  materials  were  altered  by  decomposition,  and.  that  the  products 
were  formed  by  recombination  of  these  on  a  different  plan.  A  third 
variety  of  chemical  change  consists,  therefore,  in  the  concurrence  of 
both  of  the  first  two  kinds  in  the  same  action.  As  several  different 
varieties  of  this  sort  of  complex  redistribution  of  material  can  be  dis- 
tinguished, a  distinct  name  is  given  to  each.  The  present  is  called  a 
double  decomposition  or  metathesis. 

When  we  encounter  other  chemical  changes,  we  shall  find  the 
extent  of  the  stride  we  have  taken  in  this  critical  analysis  of  a 
few  examples.  It  will  then  appear  that  the  chemical  changes  of 
matter  are  not  nearly  so  various  as  we  might  have  anticipated.  In 
fact,  there  are  few  changes  which  cannot  be  placed  in  one  of  the 
above  categories.  Those  that  cannot  be  so  placed  belong  to  a  fourth 
sort  of  change  which,  for  the  sake  of  completeness,  may  now  be 
mentioned. 

It  occasionally  happens,  especially  in  the  case  of  compounds  of 
carbon  (see  Urea),  that  one  single  kind  of  material  turns  into  another 
single  kind  of  material.  Nothing  is  added  and  nothing  removed,  yet 
the  new  substance  has  different  properties  in  every  respect  from  the 
old.  Most  of  those  substances  whose  transformation  is  definitely 
assigned  to  this  class  contain  several  constituents,  but  a  rough  notion 
of  this  sort  of  chemical  change  may  be  obtained  by  considering  the 
two  forms  of  phosphorus.  One  of  them  is  pale  yellow  in  appearance, 
easily  melted,  and  very  easily  *  ignited ;  the  other  is  red,  does  not  melt 
on  being  heated,  and  is  difficult  to  ignite.  The  latter  is  made  from 
the  former  by  heating  it  continuously  for  some  hours  in  a  closed  ves- 
sel at  about  300°.  As  no  material  is  taken  up,  the  weight  of  the  sub- 
stance is  unchanged,  and  yet,  when  the  vessel  is  opened,  the  common 
phosphorus  is  found  to  have  turned  into  the  red  variety.  As  the 
foregoing  paragraphs  show,  we  have  good  reason  to  believe  that  the 
properties  of  a  substance  are  intimately  related  to  its  constituents. 
Carrying  out  this  idea,  the  hypothesis  (or  suggestion)  has  been  made 
that,  since  here  also  the  properties  change,  there  must  be  some  read- 
justment of  the  material,  even  in  cases  like  this.  Hence,  we  designate 
changes  of  this  kind  internal  rearrangements. 


16  INORGANIC  CHEMISTRY 

In  the  completeness  of  the  transformation,  and  in  the  fact  that  only  one  origi- 
nal substance  and  one  product  are  required,  the  physical  changes  of  the  nature  of 
melting  a  solid  and  vaporizing  a  liquid  resemble  the  fourth  variety  of  chemical 
change.  Where,  then,  is  the  line  between  chemistry  and  physics  to  be  drawn?  It 
is  in  this  fourth  group  only  that  the  difficulty  is  encountered.  All  agree  that 
warm  solid  phosphorus  is  chemically  identical  with  cold  phosphorus.  Nearly  all 
scientific  men  at  present  assign  the  study  of  the  melting  and  the  vaporization  of 
phosphorus  and  other  substances  to  physics.  Some  chemists  consider  the  solu- 
tion of  phosphorus  in  carbon  disulphide  or  some  other  solvent  (it  is  practically 
insoluble  in  water),  although  the  material  is  recovered  unchanged  by  evaporation 
of  the  liquid,  as  a  chemical  change.  But  the  great  majority  regard  this  as  physi- 
cal also.  If  not  chemically  different,  the  solid,  liquid,  gaseous,  and  dissolved  forms 
of  a  substance  must  be  classed  as  mere  physical  states  of  aggregation.  On  the 
other  hand,  red  phosphorus  is  held  by  most  chemists  to  be  chemically  different 
from  yellow  phosphorus.  Many  kinds  of  matter  show  as  much  variety  in  form  as 
phosphorus,  and  some  show  more. 

The  Second  Characteristic  of  Chemical  Phenomena.  —  We 

are  now  able  to  make  another  generalization  (or  condensed  statement 
of  fact)  :  In  chemical  phenomena  substances  enter  into  combination, 
come  out  of  combination,  or  change  their  associates  in  combination  or 
their  state  in  the  compound.  In  other  words,  the  material  changes  its 
composition  or  its  constitution. 

Summary.  —  Thus  far,  we  have  learned  that  chemistry  deals  with 
the  changes  in  composition  and  constitution  which  substances  undergo, 
and  with  the  alteration  in  properties  which  accompanies  and  gives 
evidence  of  those  changes. 

Exercises.  —  1.  Take  one  by  one  the  words  or  phrases  printed  in 
black  type  and  the  titles  of  the  sections  in  this  chapter,  and  endeavor 
to  recollect  what  you  have  read  about  each.  In  each  case  try,  (a)  to 
recall  the  meaning  and  to  state  it  in  your  own  words  ;  (b)  to  recall  the 
facts  associated  with,  and  the  reasoning  which  led  up  to  the  point  in 
question ;  (c)  to  recall  examples  illustrating  the  conception  and  to 
apply  the  conception  in  detail  to  each  example.  Whenever  memory 
fails  to  give  a  perfectly  clear  report  of  the  matter  in  hand,  the  text 
must  be  read  and  reread  until  the  essential  point  can  be  repeated  from 
memory. 

Use  the  same  method  in  all  future  chapters.  A  useful  practice  is 
to  employ  a  pencil  as  you  read  and  to  underline  systematically  all  the 
important  facts  and  statements,  and  then  to  go  back  and  apply  to  each 
marked  place  the  process  described  above. 


CHAPTEE  II 
INTRODUCTORY  II 

IF  we  now  return  to  the  three  illustrations  of  chemical  phenomena 
which  we  have  been  studying  (pp.  10-14),  we  shall  find  a  new  question 
arising  naturally  out  of  them.  This  is,  whether  the  mass  of  the  ma- 
terials is  altered,  as  are  the  other  attributes,  in  these  chemical  changes. 

Third  Characteristic  of  Chemical  Phenomena  :  Conserva- 
tion of  Mass.  —  The'  most  painstaking  chemical  work  seems  to  show 
that,  if  all  the  substances  concerned  in  the  chemical  change  are 
weighed  before  and  after  the  change,  there  is  no  evidence  of  any 
alteration  in  the  quantity  of  matter.  The  two  weights,  representing 
the  sums  of  the  constituents  and  of  the  products  respectively,  are, 
indeed,  never  absolutely  identical,  but  the  more  careful  the  work  and 
the  more  delicate  the  instrument  used  in  weighing,  the  more  nearly 
do  the  values  approach  identity.  We  are  able  to  state,  therefore,  as 
a  third  characteristic  of  all  chemical  phenomena,  that  the  mass  of  a 
system  is  not  affected  by  any  chemical  change  within  the  system. 

This  statement  simply  means  that  the  great  law  of  the  conserva- 
tion of  mass  holds  true  in  chemistry  as  it  does  in  physics.  Chemical 
changes,  thoroughgoing  as  they  are  in  respect  to  all  other  properties, 
do  not  affect  the  mass ;  an  element  carries  with  it  its  weight,  entirely 
unchanged,  through  the  most  complicated  chemical  transformations. 
This  is  the  only  attribute  which  persists. 

A  law,  as  we  have  seen  (p.  7),  is  a  condensed  statement  describing  some  con- 
stant mode  of  behavior.  It  is  simply  a  summary  of  our  experience.  As  such,  it 
is  subject  to  modification  when  a  fact  is  discovered  with  which  it  conflicts.  Thus, 
it  is  perfectly  possible  that  we  may  yet  find  cases  of  demonstrable  changes  in 
weight  accompanying  other  physical  or  chemical  changes  in  a  limited  system. 
Indeed,  it  has  more  than  once  been  alleged  that  such  changes  have  been  observed. 
It  used  to  be  a  law  that  the  earth  was  flat.  It  is  now  more  correct  to  say  that  a 
limited  area  of  perfectly  level  ground  is  very  nearly  flat. 

It  will  be  observed  that  the  phrasing  of  the  above  law  carefully  limits  its 
scope  to  amounts  of  matter  such  as  are  dealt  with  in  laboratory  experience.  We 
have  no  evidence  on  which  to  make  any  statements  about  the  mass  of  matter  in 
more  extensive  chemical  changes.  A  common  form  of  the  law,  to  the  effect  that 

17 


18 


INORGANIC   CHEMISTRY 


"  the  mass  of  matter  in  the  universe  is  unchangeable  in  amount,"  is  not  a  law  at 
all,  in  the  only  sense  in  which  the  word  is  used  in  science.  It  is  a  statement  in 
regard  to  supposed  facts  which  are  almost  entirely  beyond  our  experience.  It  is, 
therefore,  a  proposition  of  a  transcendental  (that  is,  transcending  experience) 
nature,  and  has  its  proper  place  in  metaphysics.  Astronomical  observation,  it  is 
true,  has  as  yet  furnished  no  evidence  of  changes  in  the  mass  of  our  own  or  other 
celestial  systems.  But,  absence  of  evidence  to  the  contrary,  especially  considering 
the  relatively  limited  scope  of  our  knowledge,  both  in  respect  to  space  and  time,  is 
far  from  being  proof  of  the  correctness  of  the  proposition. 

Superficial  observation,  as  of  a  growing  tree,  might  seem  to  give 
evidence  of  the  very  opposite  of  conservation  of  matter.  But  here 

the  carbon  dioxide  gas  in 
the  air,  the  most  impor- 
tant source  of  nourishment 
for  plants,  is  overlooked. 
Similarly  the  gradual  dis- 
appearance of  a  candle  by 
combustion  seems  to  illus- 
trate the  destruction  of 
matter.  But  if  we  insert 
sticks  of  sodium  hydroxide 
in  a  U-tube  (Fig.  10)  to 
catch  the  gases  which  rise 
through  the  flame,  we  find 
that  the  gases  weigh  even 
more  than  the  part  of  the 
candle  which  has  been  sac- 
rificed in  making  them. 
When  we  take  account  of 
the  weight  of  the  oxygen 
obtained  from  the  air  which 
sustains  the  combustion, 
we  find  that  there  is  really 
neither  loss  nor  gain  in 

weight.  If  we  carry  out  chemical  changes  in  closed  vessels  (Fig.  11), 
which  permit  neither  escape  nor  access  of  material,  we  find  that  the 
weight  does  not  alter. 

One  way  of  stating  the  difference  between  chemistry  and  physics  is  to  say  that 
changes  in  which  both  the  mass  and  the  identity  of  the  substance  are  conserved 
belong  to  the  latter,  while  those  in  which  the  mass  alone  is  conserved  belong  to 
chemistry. 


FIG.  10. 


INTRODUCTORY  II 


19 


Physical  Concomitants  of  Change  in  Composition.  —  Study 
of  the  three'  typical  chemical  changes  described  in  the  last  chapter 
may  now  be  resumed,  in  order  to  see  whether  anything  further  of  a 
general  nature  is  characteristic  of  such  phenomena.  We  recall  at 
once  that  a  prominent  feature  of  the  union  of  iron  and  sulphur  was 
the  heat  which,  as  shown  by  the  glow 
spreading  through  the  mass,  seemed  to 
be  developed  after  the  action  was  once 
started.  It  is  found  that  many  chemi- 
cal changes  are  like  this  one,  in  exhibit- 
ing simultaneously  the  production  of 
very  perceptible  amounts  of  heat.  On 
the  other  hand,  the  decomposition  of 
mercuric  oxide,  as  was  pointed  out 
(p.  13),  owed  its  continuance  to  the  per- 
sistent application  of  heat,  and  ceased 
so  soon  as  the  source  of  heat  was  with- 
drawn. Here,  apparently,  heat  was 
consumed  during  the  progress  of  the 
change,  and  the  chemical  action  was 
limited  by  the  amount  of  heat  supplied. 
The  production  or  consumption  of  heat 
may,  therefore,  be  a  feature  of  chemical 
change. 

In  the  iron  and  sulphur  case,  as  in  other  chemical  actions  where 
the  heat  developed  is  great,  light  also  was  given  out.  In  the  last  of 
the  three  actions,  on  the  other  hand,  we  obtained  a  substance  (silver 
chloride),  which  may  be  kept  for  any  length  of  time  in  the  dark,  but, 
by  the  action  of  sunlight  is  broken  up  into  its  constituents  (p.  14).  It 
would  appear,  therefore,  that  light  may  be  given  out  or  used  in  con- 
nection with  chemical  change.  Noting  these  facts  stimulates  us  to 
look  for  other  similar  concomitants  of  changes  in  composition. 

If  we  dip  two  wires  from  a  battery  or  dynamo  into  a  solution  of 
nitrate  of  silver  (Fig.  12),  such  as  was  used  in  the  third  experiment, 
we  observe  the  instant  production  of  a  coating  of  silver  on  the  nega- 
tive wire.  By  preparing  the  solution  properly  and  allowing  the 
electricity  to  flow  through  it  for  a  sufficient  length  of  time,  all  of 
the  compound  can  be  decomposed  and  all  its  silver  deposited.  It  is 
needless  to  say  that  this  release  of  the  silver  from  chemical  combina- 
tion and  liberation  of  the  metal  at  the  electrode,  goes  on  only  so  long 


FIG.  11. 


20 


INORGANIC   CHEMISTRY 


FIG.  12. 


as  the  current  of  electricity  is  employed,  and  that  electricity  is  con- 
sumed in  the  process.     Very  many  substances  can  be  decomposed  in 

this  way. 

The  inverse  of  this  is 
likewise  familiar.  If  we 
place  in  dilute  sulphuric 
acid  a  stick  of  the  metal 
zinc,  we  find  that  a  gas 
is  given  off  rapidly  (Fig. 
13),  that  the  zinc  gradu- 
ally dissolves,  and  that  a 
large  amount  of  heat  is 
developed.  Under  favor- 
able circumstances,  the 
liquid  may  even  rise  spon- 
taneously to  the  boiling- 
point.  This  form  of  the 
action  produces  heat.  If, 

however,  we  attach  the  same  stick  of  zinc  to  a  copper  wire,  and, 
having  provided  a  plate  of  platinum  also  connected  with  a  wire, 
immerse  the  two  simultaneously  in  the  acid 
(Fig.  14),  then  a  galvanometer,  with  which  the 
wires  are  connected,  shows  at  once  the  passage 
of  a  current  of  electricity  round  the  circuit. 
Exactly  the  same  chemical  change  goes  on  as 
before.  The  sole  difference  is  that  the  gas 
appears  to  arise  from  the  surface  of  the  plati- 
num. It  is  easy  to  show,  however,  that  the  plat- 
inum by  itself  is  not  acted  upon  by  dilute  acids, 
and,  in  this  case,  undergoes  no  change  whatever ; 
it  serves  simply  as  a  suitable  conductor  for  the 
electricity.  Here,  then,  in  place  of  the  heat 
which  the  first  plan  produced,  we  get  electricity. 
The  arrangement  is,  in  fact,  a  battery,  for  a  bat- 
tery is  a  system  in  which  a  chemical  action  which 
would  otherwise  give  heat  furnishes  electricity  in- 
stead. Thus,  electricity  may  be  consumed  or  pro- 
duced in  connection  with  a  change  in  composition. 

Even  violent  rubbing  in  a  mortar,  in  the  case  of  some  substances, 
can  effect  an  appreciable  amount  of  decomposition  in  a  few  minutes. 


~S'* 


FIG.  13. 


INTRODUCTORY   II 


21 


In  this  way  silver  chloride  can  be  separated  into  silver  and  chlorine, 
just  as  by  light.  It  is  the  mechanical  energy  which  is  the  agent,  and 
part  of  it  is  consumed  in  producing  the  change,  and  only  the  balance 
appears  as  heat.  Conversely,  the  production  of  mechanical  energy,  as 


FIG.  14. 


the  result  of  chemical  change,  is  seen  in  the  behavior  of  explosives 
and  in  the  working  of  our  muscles.  Thus,  mechanical  energy  may  be 
used  up  or  produced  in  chemical  changes. 

Summing  our  experience  up,  we  may  state  that  no  change  in  com- 
position occurs  without  some  concomitant,  such  as  the  production  or 
consumption  of  heat,  light,  electricity,  or,  in  some  cases,  mechanical 
energy. 

Classification  of  the  Concomitants  of  Change  in  Composi- 
tion: Energy.  —  The  problem  of  classifying  (i.e.,  placing  in  a  suitable 
category)  things  like  heat,  light,  and  electricity  has  occupied  much 
attention.  They  do  not  possess  mass.  In  all  changes  in  composition, 
one  of  these  natural  concomitants  is  given  out  or  absorbed,  sometimes 
in  great  amount,  yet,  in  none  is  any  alteration  in  weight  observed. 
Nor  may  these  concomitants  be  overlooked,  simply.  A  conception  is 
a  real  thing ;  a  religious  belief  may  be  most  real  and  potent.  There 
are  many  things  which  are  real,  although  they  are  not  affected  by 
gravitation.  In  the  present  instance  we  reason  as  follows  : 

A  brick  in  motion  is  different  from  a  brick  at  rest.  The  former 
can  do  some  things  that  the  latter  cannot.  Furthermore,  we  can 


22 


INORGANIC   CHEMISTRY 


easily  make  a  distinction  in  our  minds.  The  brick  can  be  deprived  of 
the  motion  and  be  endowed  with  it  again.  Thus,  we  can  get  the  idea 
of  motion  as  a  separate  conception.  Similarly,  we  observe  that  a  piece 
of  iron  behaves  differently  when  hot,  and  when  cold,  when  bearing  a 
current  of  electricity,  and  when  bearing  none.  We  conceive  then 
of  the  brick  or  the  iron  as  having  a  certain  amount  and  kind  of  matter 
which  is  unalterable,  and  as  having  motion,  heat,  or  electricity  added 
to  this  or  removed.  Thus,  we  describe  our  observations  by  using  two 
categories,  one  of  which  includes  the  various  kinds  of  matter,  and 
the  other,  various  things  whose  association  with  it  seems  to  be  invari- 
able and  is  often  so  conspicuous. 

At  first  sight,   these   concomitants  of   matter   seem  to  be  quite 
disparate.     But  a  relation  between  them  can  be  found.     If  the  heat 

of  a  Bunsen  flame  or  of  the  sun  is  brought 
under  a  hot-air  motor  (Fig.  15)  violent 
motion  results.  Again,  if  the  motor  is 
connected  with  a  dynamo,  electricity  may 
be  generated.  Still  again,  if  the  current 
flows  through  an  incandescent  lamp,  heat 
and  light  are  evolved.  Conversely,  when 
motion  is  impeded  by  a  brake,  heat  appears. 
When  a  current  of  electricity  is  run 
through  the  dynamo,  motion  results.  But 
the  most  significant  facts  are  still  to  be  men- 
tioned. The  heat  absorbed  by  the  motor 
is  found  to  be  greater  when  the  machine 
is  permitted  to  move  and  do  work,  than 
when  it  is  not.  Thus,  it  is  found  that 
when  work  is  done  some  heat  disappears, 
and  is,  in  fact,  transformed  into  work. 

Similarly,  when  the  poles  of  the  dynamo  are  properly  connected 
and  electricity  is  being  produced,  and  only  then,  motion  is  used 
up.  This  is  shown  by  the  effort  required  to  turn  the  armature 
under  these  circumstances,  and  the  ease  with  which  it  is  turned 
when  the  circuit  is  open.  So,  with  a  conductor  like  the  filament 
in  the  lamp,  unless  it  offers  resistance  to  the  current  and  destroys 
a  sufficient  amount  of  electricity,  it  gives  out  neither  light  nor 
heat.  Finally,  motion  gives  no  heat  unless  the  brake  is  set,  and 
effort  is  then  demanded  to  maintain  the  motion.  These  experiences 
lead  us  to  believe  that  we  have  here  a  set  of  things  which  are  funda- 


Fio.  15. 


INTRODUCTORY   II  23 

mentally  of  the  same  kind,  for  each  form  can  be  made  from  any  of 
the  others.  We  have,  therefore,  invented  the  conception  of  a  single 
thing  of  which  heat,  light,  electricity,  and  motion  are  forms,  and  to  it 
we  give  the  name  energy :  energy  is  work  and  every  other  thing  which 
can  arise  from  work  and  be  converted  into  -work  (Ostwald). 

Closer  study  shows  that  equal  amounts  of  electrical  or  mechanical 
energy  always  produce  equal  amounts  of  heat.  There  is  never  ob- 
served any  loss  in  any  of  the  transformations  of  energy  any  more 
than  in  the  transformations  of  matter.  Hence,  J.  R.  Mayer  (1842), 
Colding  (1843),  and  Helmholtz  (1847)  were  led  independently  to  the 
conclusion  that  in  a  limited  system  no  gain  or  loss  of  energy  is  ever 
observed.  This  brief  statement  of  the  results  of  many  experiments 
is  called  the  law  of  the  conservation  of  energy. 

A  current  form  of  this  law,  namely,  that  "the  total  amount  of  energy  in  the 
universe  is  a  constant  quantity,"  is  open  to  the  same  objection  as  the  correspond- 
ingly flamboyant  form  of  the  law  of  conservation  of  mass  criticized  above.  It  has 
a  more  effective  sound  than  the  one  we  have  given.  Unfortunately,  it  is  not  only 
immensely  in  excess  of  any  statement  that  present  results  of  scientific  work  can 
justify,  but  is  probably  far  beyond  the  limits  of  possible  scientific  observation. 
Scientific  statements  of  fact  can  never  err  by  being  too  conservative. 

Matter  and  Energy  as  Concepts,  and  Definitions  of  the  Lat- 
ter. —  The  foregoing  paragraphs  about  energy  bring  up  the  question  of  its  relation 
to  matter.  This  relation  can  be  made  clear  only  by  a  somewhat  elaborate  dis- 
cussion of  our  fundamental  conceptions. 

The  only  real,  first-hand  knowledge  which  we  possess,  is  that  of  our  states  of 
consciousness.  All  else,  consisting,  for  example,  of  the  way  in  which  we  inter- 
pret and  describe  our  experience,  is  constructed  out  of  our  heads,  so  to  speak. 
Now,  we  become  aware  of  certain  things  which  we  call  sensations,  and  seek  to  con- 
struct a  mode  of  correlating  and  describing  them.  Our  universal  habit  is  to  speak 
as  if  they  were  produced  by  something  outside  our  minds,  and  so  we  begin  the 
manufacture  of  an  external  universe.  In  course  of  doing  this,  we  encounter 
some  things  which  seem  to  occupy  no  particular  space,  which  move  from  object 
to  object,  and  possess  no  weight.  One  of  these  affects  our  eye,  or  a  piece  of 
chloride  of  silver,  for  example,  yet  escapes  touch,  and  passes  through  glass  as 
easily  as  through  a  vacuum.  After  consideration  of  our  experience  with  this  sort 
of  thing,  some  of  which  has  been  detailed  above,  we  decide  that  we  shall  posit  the 
existence  of  energy. 

Other  things  we  encounter  which  appeal  to  the  sense  of  touch  and  seem  to 
possess  more  definitely  located  qualities,  including  weight.  Another  conception 
is  needed  to  account  for  these  ;  so  we  establish  the  category  of  matter. 

Thus,  we  make  shift  to  describe  our  sensations  by  the  help  of  these  two  con- 
structs, much  as  in  analytical  geometry  we  describe  the  location  of  a  point  by 
means  of  two  coordinates.  Energy  and  matter  are,  therefore,  products  of  thought 
and  not,  primarily,  objective  realities.  In  chemistry,  however,  we  always  speak 


24  INORGANIC   CHEMISTRY 

of  them  objectively.  Historically,  the  order  in  which  these  two  concepts  were 
named  and  defined  was  the  opposite  of  that  in  which  they  stand  above.  Yet 
attempts  to  organize  a  conception,  corresponding  to  energy,  in  response  to  a  need 
of  which  thinkers  were  conscious,  were  not  wanting  before  the  nineteenth  century 
opened.  To  go  no  further  back  than  the  days  of  phlogiston,  we  can  easily  per- 
ceive a  certain  resemblance  between  this  concept  and  that  of  energy.  The  idea 
that  heat  was  an  "  imponderable  "  had  its  origin  much  earlier,  and  shows  the  ex- 
istence of  the  same  effort  to  find  a  second  fundamental  conception  different  from 
matter. 

There  is  much  confusion  of  thought  in  many  of  the  current  definitions  of 
energy.  For  example,  it  is  often  said  to  be  "  that  which  causes  change  in  matter." 
This  definition  is  not  easy  to  bring  into  harmony  with  common  experience  in 
chemistry.  Thus,  when  heat  is  applied  to  mercuric  oxide,  the  change  follows. 
But  with  iron  and  sulphur,  the  union  of  the  two  substances  is  a  condition  antece- 
dent to  the  evolution  of  heat.  It  is  as  often  true  that  change  in  matter  causes  the 
manifestation  of  energy  as  the  reverse.  Matter  and  energy  are  on  the  same  plane. 
They  are  conceptions  used  jointly  in  describing  what  we  observe.  Neither  is 
secondary  to  the  other.  We  do  not  consider  any  particular  one  of  the  coordinates 
in  geometry  as  secondary  to  the  other,  or  as  being  affected  by  the  other. 

The  recent  theory  of  chemical  potential  and  of  the  factors  of  energy  (q.v.)  seeks 
once  more  to  ascribe  the  tendency  to  change  (physical  or  chemical)  in  matter  to 
the  state  of  the  energy  associated  with  it.  It  is,  therefore,  incidentally,  so  con- 
structed as  to  favor  the  definition  just  given. 

The  definition  that  «'  matter  is  the  vehicle  of  energy  "  is  obviously  just  as  diffi- 
cult to  harmonize  with  the  above  mode  of  deriving  the  two  conceptions.  One  axis 
is  not  spoken  of  as  the  vehicle  of  the  other  in  geometry. 

The  innate  desire  to  reduce  our  distinct  categories  to  the  smallest  possible 
number  may  be  seen  in  the  history  of  this  subject.  The  ancients  sought  the 
amalgamation  of  the  two  by  regarding  heat  and  light  as  imponderable  forms  of 
matter.  In  some  quarters  it  is  now  believed  that  the  one  conception  of  energy  is 
sufficient,  and  that  matter  may  be  put  into  the  same  category  as  being,  for  example, 
composed  of  minute  particles  of  electricity  (electrons). 

The  conception  of  ether  was  devised  because  those  of  matter  and  energy  did 
not  suffice  for  the  description  of  all  the  phenomena  of  light.  It  is  on  the  same 
plane  with  matter  and  energy.  Lord  Kelvin's  effort  to  reduce  the  three  catego- 
ries to  two  (energy  and  ether)  by  assigning  the  rdle  of  matter  to  vortices  in  the 
ether  is  familiar  to  students  of  physics. 

Application  of  the  Conception  of  Energy  in  Chemistry.  —  At 

first  sight  it  looks  as  if  the  statement  that  energy  is  conserved  is  not 
applicable  in  chemistry.  Heat  and  electricity,  for  example,  seem  to 
be  produced  and  consumed,  in  connection  with  changes  in  composition, 
in  a  mysterious  manner.  We  trace  light  in  an  incandescent  lamp 
back  to  the  electricity,  and  this  in  turn  to  the  mechanical  energy,  and 
this  again  to  the  heat  in  the  engine.  But  what  form  of  energy  gave 
the  heat  developed  by  the  combustion  of  the  coal  under  the  boiler,  or 


INTRODUCTORY   II  25 

by  the  union  of  iron  and  sulphur  in  our  first  experiment  ?  Since  we 
do  not  perceive  any  electricity,  light,  heat,  or  motion,  in  the  original 
materials,  and  yet  wish  to  create  an  harmonious  system,  we  are  bound 
to  conceive  of  the  iron  and  the  sulphur,  and  the  coal  and  the  air,  as 
containing  another  form  of  energy,  which  we  call  chemical  or  internal 
energy.  Similarly,  when  heat  is  used  up  in  decomposing  mercuric 
oxide,  or  light  in  decomposing  silver  chloride,  we  regard  the  energy  as 
being  stored  in  the  products  of  decomposition  in  the  form  of  chemical 
energy. 

The  Actual  Quantities  of  Different  Kinds  of  Energy  which 
may  be  Obtained  from  a  Fixed  Amount  of  One  Kind.  —  It  will 
render  all  the  above  clearer  if  we  give  some  numerical  illustrations :  A  kilogram 
of  water  after  falling  (in  a  vacuum)  428  meters  (about  one-fifth  of  a  mile),  under 
gravity,  possesses  428  kilogram-meters  of  mechanical  (kinetic)  energy.  When  the 
motion  is  arrested,  the  energy  of  motion  is  transformed  into  heat  and  raises  the 
water  one  degree  centigrade  in  temperature.  We  describe  this  amount  of  heat 
as  1000  calories  (small)  ;  that  required  to  warm  one  gram  of  water  one  degree 
(between  0°  and  100°)  being  called  one  calorie.  A  kilogram  of  any  other  falling 
material  would  give  the  same  amount  of  heat  (1000  cal.),  although,  of  course, 
if  its  specific  heat  were  smaller  than  that  of  water,  the  temperature  to  which  it 
would  be  raised  would  be  higher,  and  vice  versa. 

Here  the  acting  force  is  the  attraction  of  gravitation,  which  is  a  special  case. 
In  absolute  units,  1  g.  falling  1  cm.  generates  energy  enough  to  do  981  ergs  of  work. 
So  that  the  thousand  grams  falling  428  meters,  generates 

1000x42,800x981=42,000,000,000  ergs  of  energy. 

The  erg  being  so  small,  we  often  use  the  joule  (  =  10,000,000  ergs).  This  amount 
is  the  same  as  4200  joules. 

Now,  any  body  of  the  same  mass,  moving  with  the  same  final  velocity,  however 
set  in  motion,  will  also  possess  the  same  energy  and  give  1000  cal.  The  final 
velocity  in  the  above  case  is  V2  gs.  =  9,164  cm.  per  second.  The  energy  of  motion 
(£  rav2)  of  one  thousand  grams  of  matter  moving  with  this  velocity  is  =  |  X  1000  x 
(9164)2  =•=  4200  joules,  as  before.  If  the  source  of  mechanical  energy  were  a  hot- 
air  motor  (or  an  engine)  of  one  horse-power,  then,  since  one  horse-power  repre- 
sents a  development  of  746  joules  per  second,  the  4200  joules  of  energy  would  be 
produced  in  about  5J  seconds  by  this  means. 

If,  instead  of  being  turned  into  heat,  all  the  energy  of  motion  had  been  con- 
verted into  electricity,  the  quantity  of  the  latter  would  have  illuminated  a  16 
candle-power  incandescent  lamp  for  84  seconds  (  =  1.4  minutes).  Such  a  lamp  re- 
quires the  value  of  50  joules  per  second  of  electricity  and,  therefore,  in  84  seconds 
uses  up  the  50  x  84  =  4200  joules  of  energy.  As  an  engine  of  I  horse-power  pro- 
duces this  amount  of  energy  every  5J  seconds,  such  an  engine,  if  none  of  the  energy 
were  lost,  could  maintain  nearly  15  lamps  of  this  kind. 

Finally,  if  the  4200  joules  of  electrical  energy  were  applied  to  decomposing 
nitrate  of  silver  in  ordinary  aqueous  solution,  it  would  liberate  6$  grams  (about 
^  oz.)  of  silver  from  combination. 


26  INORGANIC   CHEMISTRY 

Considerations  Connected  with  Chemical  Energy:  Free 
Energy.  —  These  conclusions  compel  us,  for  the  sake  of  consistency,  to 
think  of  all  our  materials  as  repositories  of  energy  as  well  as  of  mat- 
ter, each  of  these  constituents  being  equally  real  and  equally  impor- 
tant. A  piece  of  the  substance  known  as  "  iron  "  must  thus  be  held  to 
contain  so  much  iron  matter  and  so  much  chemical  energy.  So  ferrous 
sulphide  contains  sulphur  matter,  iron  matter,  and  chemical  energy. 
Thus,  by  a  substance  we  mean  a  distinct  species  of  matter,  simple  or 
compound,  with  its  appropriate  proportion  of  chemical  energy.  Dur- 
ing the  progress  of  a  chemical  change  like  the  union  of  iron  and 
sulphur,  the  chemical  energy  of  the  system  diminishes  and  heat  is 
liberated,  or,  when  arrangements  are  made  for  utilizing  the  energy, 
work  of  some  kind  is  done. 


The  energy  which  becomes  available  as  the  result  of  a  chemical  action,  and 
is  free  to  be  converted,  say,  into  electrical  energy,  is  called  the  free  energy  of 
the  action.  Now,  it  must  be  noted  that  the  free  energy,  measured  by  work  done, 
is  not,  in  general,  the  equivalent  of  the  heat  developed  by  the  natural  progress  of 
the  change.  Often  the  amounts  are  nearly  equivalent,  although  never  absolutely 
so.  But  frequently  they  are  very  different.  When  the  free  energy  available  for 
conversion  into  work  is  greater  in  amount  than  the  heat  of  the  reaction,  as  it  often 
is,  the  difference  is  taken  up  from  the  heat  of  the  surroundings  during  the  progress 
of  the  change,  and  the  vessels  and  objects  in  contact  with  the  interacting  bodies 
become  colder.  Thus  phosphonium  chloride  (g.  t>.)  decomposes  spontaneously  into 
two  gases,  phosphine  and  hydrogen  chloride,  and  ammonium  carbonate  gives  off 
ammonia  gas,  while  heat  is  absorbed  in  both  cases.  Work  can  be  done  by  both 
these  actions,  although,  so  far  as  heat  is  concerned,  not  only  is  none  of  this  form 
of  energy  liberated,  but  a  certain  amount  of  it  is  absorbed.  Conversely,  when  the 
heat  of  reaction  is  greater  than  the  equivalent  of  the  free  energy,  then,  along  with 
the  energy  which  could  be  used  to  do  work  (for  example,  by  employing  the  action 
as  a  source  of  electricity),  a  certain  amount  of  heat  which  cannot  be  transformed 
into  work  will  be  given  to  the  surroundings.  It  thus  appears  that  the  substances 
which  we  handle  are  not  only  repositories  of  energy,  but,  when  brought  together, 
also  play  the  part  of  machines  for  transforming  energy  which  they  take  from  or 
give  to  the  surroundings. 


The  Fourth  Characteristic  of  Chemical  Phenomena.  —  In  the 

course  of  this  discussion  it  has  become  clear  that  it  is  a  characteristic 
of  a  chemical  phenomenon  that,  besides  a  change  in  the  state  of  the 
'matter -,  there  is  always  an  alteration  in  the  amount  of  chemical  energy 
in  the  system.  This  alteration  involves  the  production  of  chemical 
energy  from,  or  the  transformation  of  chemical  energy  into,  some  other 
form  of  energy. 


INTRODUCTORY  II  27 

The  energy  liberated  in  a  chemical  action  appears  most  commonly 
in  the  form  of  heat.  Changes  which,  like  the  union  of  iron  and 
sulphur  (p.  11),  are  accompanied  by  the  liberation  of  heat  are  called 
exothermal  actions.  Those  in  connection  with  which  heat  is  absorbed, 
like  the  decomposition  of  mercuric  oxide  (p.  12)  or  of  phosphonium 
chloride,  are  known  as  endothermal. 

It  should  be  noted  here  that  neither  the  production  nor  the  ab- 
sorption of  heat  is  an  exclusive  mark  of  chemical  change.  Physical 
changes  are  all  likewise  accompanied  by  the  liberation  or  consumption 
of  energy.  Thus,  water,  in  evaporating,  absorbs  heat,  and  liquids  on 
solidifying,  or  often  even  when  simply  mixed  with  other  liquids,  give 
out  heat. 

The  absorption  or  liberation  of  energy  accompanying  a  chemical 
transformation  of  matter  is  often,  of  the  two,  the  more  important  fea- 
ture. We  do  not  burn  coal  in  order  to  manufacture  carbon  dioxide 
gas.  We  are  glad  to  get  rid  of  the  material  product  through  the  chim- 
ney. It  is  the  heat  we  want.  We  do  not  employ  zinc  in  batteries 
with  the  object  of  making  zinc  chloride  or  zinc  sulphate.  So  we  use 
the  electrical  energy,  and  throw  the  material  products  away  when 
we  refill  the  battery  jars.  It  is  the  same  with  burning  illuminating- 
gas  or  magnesium  powder  when  we  want  light,  and  with  eating  food, 
which  we  do,  chiefly,  to  get  energy  to  sustain  our  activity.  We  do 
not  run  electricity  for  hours  into  a  storage  battery  in  order  to  make 
a  particular  compound  (lead  dioxide,  for  example),  but  in  order  to 
save  and  store  the  energy  for  future  use.  In  industry  and  life,  fully 
half  the  total  amount  of  chemical  change  involved,  is  set  in  motion 
by  us,  solely  on  account  of  the  energy  changes  it  involves. 

As  will  be  seen  in  the  following  section,  observation  of  the 
amount  of  the  energy  absorbed  or  liberated  in  chemical  changes  is 
also  of  the  greatest  importance  in  the  scientific  study  of  chemical 
phenomena. 

Chemical  Activity.  —  Other  things  being  equal,  actions  in  which 
there  is  a  relatively  large  loss  of  chemical  energy,  and,  therefore, 
usually,  a  considerable  liberation  of  heat  or  electrical  energy,  proceed 
rapidly ;  that  is  to  say,  in  them  a  large  proportion  of  the  material  is 
changed  in  the  unit  of  time.  Those  in  which  less  free  energy  is 
transformed  proceed,  in  general,  more  slowly.  The  speed  of  the 
chemical  change,  and  the  quantity  of  energy  available  because  of  it, 
are  closely  related.  Now,  we  are  accustomed  to  speak  of  materials 


28  INORGANIC  CHEMISTRY 

which,  like  iron  and  sulphur,  interact  rapidly  and  with  liberation  of 
much  energy  as  "  chemically  active."  Thus,  relative  chemical  activity 
may  be  estimated : 

1.  By  observing  the  speed  of  a  change   (see    Speed  of  chemical 
actions),  or,  in  many  cases, 

2.  By  measuring  the  heat  developed  in  the  course  of  the  action 
(see  Thermochemistry),  or, 

3.  By  ascertaining    the    electromotive   force   of    the   current    the 
change  gives,  when  arranged  in  the  form  of  a  battery-cell  (see  Electro- 
motive chemistry). 

These  different  methods  will  be  discussed  in  later  sections.  It 
should  be  noted  here,  however,  that  the  speed  of  a  given  action  may 
be  enormously  affected  by  conditions  (see,  for  example,  Catalysis), 
and  that,  therefore,  great  caution  is  required  in  inferring  relative 
activities  from  observed  differences  in  the  speeds  of  several  actions. 
Thermal  measurements  are  also  often  misleading.  This  is  evident 
from  the  fact  that  an  action  may  be  able  to  do  work,  even  although 
heat  energy  is  absorbed  during  its  progress  (p.  26).  The  electrical 
method  of  measuring  the  free  energy,  and  therefore  the  true  affinity, 
is  in  general  the  most  trustworthy. 

It  is  evident  that  the  chemical  activity  of  a  given  substance  will 
not  be  the  same  towards  all  others.  Thus,  iron  unites  much  more 
vigorously  with  chlorine  than  with  sulphur,  and,  with  identical  amounts 
of  iron,  more  heat  is  liberated  in  the  former  case  than  in  the  latter. 
With  silver,  sodium,  and  many  other  substances,  iron  does  not  unite 
at  all.  One  of  the  tasks  of  the  chemist  is  to  make  such  comparisons 
as  this  (see  Specific  chemical  properties,  p.  67).  Evidently,  the  sub- 
stances containing  the  most  chemical  energy  will  be  in  general  the 
most  active. 


The  "Cause"  of  Chemical  Activity.  —  The  reader  will  un- 
doubtedly be  inclined  to  inquire  whether  we  can  assign  any  cause  for 
the  tendency  which  substances  have  to  undergo  chemical  change. 
Why  do  iron  and  sulphur  unite  to  form  ferrous  sulphide,  while  other 
pairs  of  elements  taken  at  random  will  frequently  be  found  to  have 
no  effect  upon  one  another  under  any  circumstances  ?  This  question 


INTRODUCTORY   II  21 

is  so  likely  to  occur  to  the  reader  that  it  should  be  dealt  with  at  once. 
The  answer  is  that  we  do  not  know.  Questions  like  this  have  to  go 
without  answer  in  all  sciences.  What  is  the  cause  of  gravitation? 

We  know  the  facts  which  are  associated  with  the  word the  fact 

that  bodies  fall  towards  the  earth,  for  example  —  but  why  they  fall 
we  are  unable  to  say.  So,  with  chemical  change,  we  can  state  all  the 
facts  we  know  about  it,  but  even  then  we  cannot  say  why  it  takes  place. 

The  words  "affinity"  and  "attraction"  are  sometimes  advanced  as  if  they 
supplied  some  explanation  of  chemical  activity.  Now,  we  have  seen  that  an  ex- 
planation in  science  (p.  10)  is  a  description  of  the  details  of  some  process,  either 
in  terms  of  known  facts,  or  by  the  use  of  some  imaginary  but  plausible  and  help- 
ful machinery.  Here  no  facts  are  known.  Even  imaginary  machinery  has  not 
yet  been  conceived  by  any  one.  So  that  these  terms  are  words  simply,  and  do 
not  meet  either  of  the  conditions  required  of  an  explanation.  They  are  names 
for  "  the  tendency  to  undergo  chemical  change,"  and  that  is  all. 

All  nouns,  such  as  table  or  book,  are  general  terms  applicable  to  many  more 
or  less  various  individuals.  Some  special  nouns  are  used  in  chemistry.  For  ex- 
ample, affinity  names  the  tendency  to  undergo  chemical  change,  and  distinguishes 
this  tendency  by  name  from  cohesion,  or  the  tendency  to  unite  physically. 
Catalysis  names  a  kind  of  chemical  change  in  which  some  specific  substance  must 
be  present,  yet  itself  undergoes  no  change.  Dissociation  names  the  kind  of  chemi- 
cal change  in  which  decomposition  occurs  with  rise  in  temperature,  and  recombi- 
nation when  the  temperature  falls.  But  none  of  these  terms,  as  such,  is  an  explana- 
tion. It  does  not  explain  the  concussion  of  two  railway  trains  to  name  it  a 
collision. 

Of  course,  if  we  have  some  genuine  explanation,  applicable  to  all  the  other 
known  cases  of  a  class,  any  newly  discovered  example  falls  heir  at  once  to  this 
explanation.  This  would  be  true  of  a  dissociation,  where  the  kinetic  theory  and 
the  law  of  mass  action  describe  the  details  of  all  such  phenomena.  Here  the  ex- 
planation lies,  not  in  the  name,  but  in  the  knowledge  we  have  of  other  instances 
of  the  same  behavior.  The  name,  of  course,  suggests  the  whole  theory,  if  such  a 
theory  exists.  But  with  affinity,  or  the  tendency  to  enter  into  chemical  action,  we 
have  no  theory  for  any  of  the  samples  of  the  class.  We  are  entirely  ignorant  as 
yet  of  the  details  of  its  mode  of  operation,  equally  so  in  every  case,  and,  in  fact, 
know  nothing  at  all  about  it  save  that  affinity  exists  and  that  we  can  measure  its 
intensity.  So  the  name  cannot  remind  us  of  any  explanation,  for  none  has  been 
suggested. 

As  words,  the  best  one  can  say  of  them  is  that  they  are  rather  unfortunately 
chosen.  Affinity  suggests  kinship,  sympathy,  or  affection.  But  the  suggestion 
that  such  human  emotions  control  the  behavior  of  iron  and  sulphur  is  too  wild 
and  too  remote  from  common  sense  to  furnish  any  assistance.  Attraction  hints 
at  some  preexisting  bond  of  a  material  kind  which  draws  the  substances  together, 
for  we  cannot  conceive  of  action  at  a  distance  without  some  intervening  medium 
of  communication.  But  we  have  no  other  evidence  of  the  existence  of  an  instan- 
taneously adjustable  harness  capable  of  drawing  materials  into  chemical  action. 
It  is  harder  to  reduce  this  idea  to  comprehensible  shape  than  to  do  without  it. 


30  INORGANIC  CHEMISTRY 

If  we  are  sometimes  inclined  to  think  that  these  are  more  than  class-words, 
and  do  suggest  some  explanation,  we  have  only  to  carry  the  same  idea  further  to 
be  landed  in  absurdity.  Using  similarly  crude  analogies,  we  might  suppose  that 
the  elements  were  guided  by  scent,  like  dogs,  or  by  sight,  like  birds,  or  by  feeling, 
like  fish,  and  so  on  ad  infinitum,  and  forget  that  the  fact  itself  was  after  all  much 
simpler  than  the  explanation.  Affinity  is  simply  a  fanciful  name  for  a  real  thing. 

"Cause"  in  Science.  —  The  word  "cause"  was  employed  in  the 
heading  of  the  last  section,  and  it  will  be  observed  that  no  cause  was 
found.  This  is  the  invariable  rule  in  physical  or  chemical  phenom- 
ena. We  know  of  no  causes,  in  the  sense  in  which  the  word  is  com- 
monly employed. 

The  word  has  only  one  definite  use  in  science.  When  we  find 
that  thorough  incorporation  of  the  three  materials  is  needed  to  secure 
good  gunpowder,  we  say  that  the  intimate  mixing  is  a  cause  of  its 
being  highly  explosive.  By  this  we  simply  mean  that  intimate  mix- 
ture is  a  necessary  antecedent  of  the  result.  A  cause  is  a  condition  or 
occurrence  which  always  precedes  another  condition  or  occurrence. 

Misuse  of  the  word  ' '  cause ' '  is  frequent.  The  law  of  gravitation  is  not  the  cause 
of  the  behavior  of  falling  bodies.  It  is  simply  a  condensed  narration  of  the  facts 
about  falling  bodies,  and  was  made  long  after  the  first  bodies  fell.  Affinity,  or 
the  tendency  to  interact  chemically,  is  the  imagined  antecedent  of  chemical  change. 
Such  causes,  if  we  call  them  causes  at  all,  are  invented -by  way  of  supplying  ante- 
cedents to  things  that  appear  to  lack  them,  that  our  sense  of  symmetry  may  be 
satisfied  withal.  They  are  occasionally  useful.  But  the  less  fiction  we  employ 
in  the  science,  the  less  will  be  the  danger  that  the  student  will  mistake  fictions  for 
facts,  or  even  fall  under  the  delusion  that  it  is  a  habit  of  science  to  spend  more 
thought  in  making  gratuitous  assumptions  than  in  ascertaining  facts. 

Of  Simple  and  Compound  Substances.  —  If  we  place  before 
a  physicist  samples  of  iron,  ferrous  sulphide,  and  sulphur,  he  will  re- 
port that  there  are  three  absolutely  distinct  substances  represented, 
because  they  show  three  different  sets  of  physical  properties.  A 
chemist,  on  the  other  hand,  while  admitting  the  accuracy  of  the  re- 
port, in  view  of  the  criterion  used  by  the  physicist,  which  indeed  he 
uses  himself  (cf.  p.  5),  will  insist  that  there  are  only  two  perfectly 
distinct  kinds  of  matter  in  the  set,  because  he  can  make  the  second  from 
matter  furnished  by  the  other  two.  The  same  sharp  contrast  in  the 
points  of  view  arises  when  mercury,  mercuric  oxide,  and  oxygen,  or 
any  similar  set  of  substances,  is  submitted  to  the  same  two  tribunals. 
In  a  sense,  chemistry  reduces  the  kinds  of  different  matter  to  a  much 
smaller  number  than  does  physics  or  any  of  the  other  sciences,  and 
so  it  is  the  final  authority  in  all  questions  involving  matter.  By  the 


INTRODUCTORY    II  31 

chemist,  dozens  of  physically  distinct  substances  are  regarded  as 
closely  related  because  they  all  can  be  made  with  iron,  or  when  de- 
composed give  it ;  hundreds  are  alike  in  that  sulphur  enters  into  their 
composition ;  thousands  are  compounds  of  oxygen.  In  fact,  the  num- 
ber of  kinds  of  matter  which  are  perfectly  distinct  in  the  strictly 
chemical  point  of  view  is  quite  limited. 

The  conception  contained  in  the  last  statement  was  not  reached 
until  centuries  of  effort  had  been  spent  in  trying  to  make  gold  out  of 
pyrite  (a  shining  yellow  mineral),  silver  out  of  lead,  and  similar 
fruitless  tasks.  The  first  to  put  our  modern  view  into  definite  lan- 
guage was  Lavoisier  in  his  Traite  de  Chimie  (1789).  His  investiga- 
tions had  already  entirely  overthrown  the  hypothesis  of  the  phlogis- 
tians  (p.  9),  and  had  been  based  upon  simple  ideas  of  combination 
or  decomposition  of  different  kinds  of  matter.  His  work  showed  that 
decomposition  had  its  limits.  Mercuric  oxide  could  be  decomposed 
into  mercury  and  oxygen,  but  no  means  was  found  of  breaking  these 
up  in  turn  and  producing  any  fresh  substances  from  them.  The  kinds 
of  matter  composing  these  simple  materials  he  named  elements.  The 
element  is  to  be  regarded  as  an  ultimate  chemical  individual  just  as 
the  substance  is  the  physical  individual.  The  definition  of  an  ele- 
ment is  therefore :  a  distinct  species  of  matter  -which  has  not  been 
shown  to  be  composite. 

The  caution  which  prompted  Lavoisier  to  use,  as  he  did,  the  words 
"  not  yet  been,"  was  justified  by  the  fact  that  several  substances,  in 
his  time  regarded  as  elementary,  were  afterwards  shown  to  be  com- 
pound. Thus,  quicklime  was  a  simple  substance  until  Davy,  in  1808, 
prepared  the  metal  calcium  and  showed  that  quicklime  was  a  com- 
pound of  this  metal  with  oxygen.  Discoveries  similar  to  this  have 
been  made  on  more  than  one  occasion  since. 

Until  recently  a  body  made  up  of  one  or  more  specific  elements  had  never 
been  found  to  yield  any  simple  substance  different  from  those  used  in  preparing 
it.  In  other  words,  one  element  had  never  been  turned  into  another.  And  this 
is  still  true  so  far  as  the  familiar  elements  are  concerned.  Experience  with  radium 
(^.i;.),  which  seems  to  decompose  and  give  helium,  however,  may  lead  to  a  revis- 
ion of  the  conception  of  an  element.  Indeed,  even  the  more  radical  hypothesis, 
that  all  the  so-called  elements  are  made  of  one  fundamental  material,  has  never 
lacked  supporters.  Perhaps  the  recent  discoveries  in  regard  to  electrons  will  lead 
to  the  conclusion,  already  formulated  by  J.  J.  Thomson,  that  the  kinds  of  simple 
matter  differ  only  in  the  number  and  arrangement  of  the  corpuscles  constituting 
their  atoms. 


32  INORGANIC  CHEMISTRY 

Element  and  Simple  Substance  Not  Equivalent  Terms.— 

We  have  seen  that  all  substances,  that  is,  physical  individuals,  must 
be  thought  of  as  containing  both  matter  and  energy  (p.  26).  It  is 
certain  that  iron,  if  it  could  be  deprived  of  the  energy  it  loses  in  com- 
bining with  sulphur,  would  be,  before  combination,  as  different  from  the 
free  metal  iron  as  is  the  compound  ferrous  sulphide  itself.  Now  the 
development  and  use  of  the  idea  of  elements  just  given,  shows  that 
by  the  word  "  element,"  applied  to  iron,  for  example,  the  chemist  refers 
only  to  the  matter  of  this  specific  variety.  It  is  this  part  alone  which 
passes  unchanged  from  the  state  of  metallic  iron  into  combination  and 
vice  versa,  or  from  one  state  of  combination  to  another.  An  element  is 
therefore  a  kind  of  matter  which  never  exists  alone.  It  is  always 
combined  with  more  or  less  chemical  energy,  and  often  with  some 
other  element  or  elements  as  well.  A  simple  substance,  on  th? 
other  hand,  like  all  substances,  has  independent  existence,  and  con 
tains  but  one  element  combined  with  a  certain  quantity  of  energy. 
A  compound  substance,  or  compound,  contains  more  than  one 
element  together  with  a  certain  amount  of  energy. 

That  this  interpretation  of  the  habit  of  thought  of  chemists  is  correct  is  shown 
by  the  fact  that  they  speak  of  both  the  yellow  and  the  red  form  as  "the  element 
phosphorus,"  and  also  describe  phosphorus  pentoxide  and  phosphine  as  "  contain- 
ing the  element  phosphorus."  The  only  thing  that  is  common  to  all  four  is,  of 
course,  matter  of  a  specific  variety,  to  which,  therefore,  the  term  "element  phos- 
phorus" must  primarily  apply.  The  individuality,  marked  by  certain  specific 
properties,  of  the  yellow  or  the  red  variety  or  of  one  of  the  compounds  as  a  sub- 
stance is  as  much  dependent  on  the  energy  it  contains  as  on  the  matter.  Hence  the 
term  substance  covers  the  energy  as  well  as  the  one  or  more  kinds  of  matter  con- 
tained in  the  material. 

Having  regard  to  this  distinction,  an  element  was  defined  above  as  a  "  variety 
of  matter,"  and  not  as  a  substance,  and  the  phraseology  used  throughout  the  parr 
graph  maintains  this  distinction.  No  mention  was  made  of  energy  in  the  defini- 
tion or  elsewhere,  for,  in  comparing  simple  substances  and  compounds  in  the 
abstract,  the  only  distinction  is  in  the  matter.  Following  each  simple  variety  of 
matter  into  and  out  of  combination,  or  from  one  state  of  combination  to  another, 
is  one  of  the  tasks  of  the  chemist.  It  is  also  his  business  to  observe  the  energy 
changes.  But  this  is  an  exercise  quite  distinct  from  the  other,  and  constitutes  a 
separate  task  in  which  different  methods  of  observation  have  to  be  employed. 

The  element  is  thus  an  abstraction,  reached  by  leaving  the  energy  out  of  COD 
sideration  and  thinking  of  the  matter  only.     It  is  invented  for  the  purpose  < 
describing  the  difference  between  the  concretes,  the  simple  substance  and  the  con. 
pound  substance.     To  use  element  and  compound  as  if  they  were  the  concrete 
things  to  be  contrasted,  involves  confusion  by  mixing  two  distinct  categories. 

As  we  have  only  one  name  for  the  element  and  the  corresponding  simple  sub- 
stance, it  is  well,  when  there  is  danger  of  ambiguity,  to  call  the  simple  substance, 
not  *'  iron,"  "  sulphur,"  or  "  the  element,"  but/ree  iron,  free  sulphur,  and  in  geL 


INTRODUCTORY  II  33 

eral  terms  the  free  element,  or  the  elementary  substance.  The  same  kinds  of  matter 
in  combination  are  then  the  element  iron  or  combined  iron,  and  the  element  sulphur 
or  combined  sulphur. 

The  chemist's  work  is  directed  wholly  by  the  thought  that  the  individual  ele- 
ment (the  matter),  after  combination,  is  still  present  in  the  compound  in  some  form 
which  is  at  least  gwcm-discrete.  The  readiness  of  the  element  to  be  released  once 
more  under  suitable  conditions  seems  to  favor  this  point  of  view.  But  with  the 
energy  in  a  substance  it  is  different.  We  cannot  easily  think  of  the  portions  of 
this,  which  came  in  with  the  several  constituents,  as  being  any  longer  attached  to 
particular  parts  of  the  compound.  Indeed,  in  some  actions  (chiefly  exothermal) 
the  free  energy  diminishes  and  a  part  escapes,  while  in  others  the  reverse  occurs. 
Whatever  energy  the  compound  contains,  it  possesses  as  a  whole,  and  a  state  of  dis- 
tribution between  the  constituents  is  seldom  taken  into  consideration.  The  specu- 
lations concerning  the  relations  of  the  atoms  in  complex  molecules,  which  have 
played  a  large  and  useful  part  in  organic  chemistry,  seem  to  constitute  the  only 
notable  divergence  from  this  point  of  view. 

These  remarks  are  introduced  because  they  help  to  demonstrate  the  logical 
necessity  of  the  chemist's  habit  of  maintaining  a  sharp  separation,  in  his  mind  at 
least,  between  questions  involving  changes  in  the  state  of  combination  of  matter 
and  questions  of  energy  (see  Formulae  and  equations,  p.  57) . 

Among  the  substances  which  we  have  been  handling,  iron,  sulphur, 
mercury,  oxygen,  and  hydrogen  are  the  free  forms  of  elements.  On 
the  other  hand,  the  substances  which  we  have  shown  to  be  composite 
are  ferrous  sulphide,  rust,  mercuric  oxide,  silver  nitrate,  and  common 
salt.  It  will  be  seen  that  by  combination  of  a  limited  number  of  ele- 
ments, two,  three,  or  four  together,  in  varying  proportions,  all  the 
known  distinct  substances  might  easily  be  accounted  for.  The  list  of 
elements  whose  individuality  has  been  established  appears  upon  another 
page  (Chap,  xii) ;  of  these  the  larger  number  are  not  frequently 
encountered.  More  than  99  per  cent  of  terrestrial  material  is  made  up 
of  eighteen  or  twenty  elements,  of  which  the  quantities  of  the  first 
eleven,  as  estimated  by  F.  W.  Clarke,  are  given  in  the  following  table : 

Oxygen 49.98  Calcium 3.51     Hydrogen 0.94 

Silicon 25.30  Magnesium    2.50    Titanium 0.30 

Aluminium 7.26  Sodium 2.28    Carbon 0.21 

Iron    5.08  Potassium 2.23  99.61 

The  evidence  of  the  spectroscope  shows  that  the  sun  and  stars  con- 
tain many  of  the  very  same  elements  as  does  the  earth. 

Some  of  the  Fundamental  Ideas  used  by  Chemists  and  the 
Corresponding  Terms.  —  To  the  chemist  it  is  above  all  important 
that  he  should  be  able  to  describe  in  unambiguous  terms  the  phenom- 
ena he  observes  and  the  inferences  he  draws. 

To  do  this  he  requires  some  fundamental  ideas  connoted  by  suit- 


34  INORGANIC   CHEMISTRY 

able  terms.  Many  of  these  ideas  and  terms  we  have  been  employing 
in  order  to  accustom  the  reader  to  their  use.  It  is  now  advisable  to 
take  them  up  more  systematically. 

Any  particular  specimen  of  matter,  such  as  a  piece  of  sulphur,  a 
portion  of  water,  a  piece  of  ferrous  sulphide,  a  fragment  of  granite,  or 
some  nitrate  of  silver  solution,  we  call  a  body.  There  are  thus  as 
many  bodies  as  there  are  discrete  portions  of  matter.  A  body  may  be 
heterogeneous,  or  made  up  of  visibly  unlike  parts,  as  granite  and  a 
mixture  of  iron  powder  and  sulphur  are ;  or  it  may  be  homogeneous, 
or  alike  in  all  parts,  as  are  pieces  of  sulphur  and  ferrous  sulphide  and 
portions  of  water  and  nitrate  of  silver  solution. 

Examination  of  these  homogeneous  bodies  shows  that  the  sulphur, 
ferrous  sulphide,  and  water  differ  from  the  nitrate  of  silver  solution  in 
having  but  one  physical  component,  while  the  last  contains  two  com- 
ponents, nitrate  of  silver  and  water,  separable  by  physical  means. 
The  last  is  like  a  mixture,  only  it  is  homogeneous.  Again,  the  first 
three  differ  amongst  themselves,  the  first  being  a  simple  body,  with 
one  chemical  constituent,  and  the  two  others  being  compounds  having 
each  two  chemical  constituents.  We  speak  of  the  components  of  a 
mixture  or  a  solution  because  the  parts  are  laid  together  and  retain  in 
the  former  case  all,  and  in  the  latter  much,  of  their  identity.  But  of  a 
compound  we  use  the  word  constituents  because  the  parts  are  built  into 
each  other  and  have  lost  their  identity. 

When  a  body,  say  a  specimen  of  sulphur,  contains  a  little  of  some 
other  physical  component,  we  speak  of  it  as  impure  sulphur.  This 
does  not  mean  that  it  contains  dirt  in  the  ordinary  sense  of  the  term. 
A  little  magnesium  chloride  is  a  common  impurity  in  table  salt  (sodium 
chloride),  and,  by  absorbing  moisture,  renders  it  more  moist  in  damp 
weather  than  it  would  otherwise  become.  "  Chemically  pure  "  means 
that  the  quantities  of  the  impurities  which  the  material  is  most  apt  to 
contain  have  been  reduced  below  the  amount  which  would  interfere 
with  the  most  exact  chemical  work  for  which  the  substance  is  com- 
monly employed.  Absolutely  pure  bodies  are  unknown. 

By  convention  we  continually  speak  of  "pure"  hydrochloric  acid, 
or  of  "pure  "  sulphuric  acid,  although  there  may  be  more  than  60  per 
cent  of  water  present  in  the  former,  and  7  per  cent  in  the  latter.  By 
this  we  mean  to  distinguish  the  former,  for  example,  from  "  commer- 
cial "  hydrochloric  acid,  which  contains  impurities  like  sulphuric  acid 
and  a  coloring  matter  in  addition  to  the  water.  The  water  is  in  fact 
disregarded,  since  it  is  assumed  to  be  present  in  all  cases. 


INTRODUCTORY  II  85 

We  distinguish  one  body,  say  one  piece  of  sulphur,  from  another 
by  its  weight,  form,  or  volume.  Each  particular  specimen  differs  from 
every  other  in  these  attributes.  These  are  general  attributes  possessed 
by  matter  and  are  used  in  chemistry  for  measuring  quantity. 

When  all  the  bodies  to  which  we  should  apply  the  name  "  sulphur  " 
are  compared,  we  find  that,  although  some  are  in  fine  powder,  and 
others  in  lumps  of  various  shapes  or  in  crystals,  and  thus  differ  in 
weight,  form,  and  volume,  they  nevertheless  have  many  qualities  in 
common.  These  qualities  we  call  specific  properties,  or  properties 
common  to  a  species.  The  material  composing  all  the  bodies  of  one 
species  we  call  a  substance.  Some  of  the  specific  properties  charac- 
terizing a  substance  and  common  to  all  specimens  of  one  species  are 
color,  odor,  crystalline  structure,  hardness,  melting-point  (temperature 
of  fusion),  solubility  in  water  or  other  solvents,  boiling-point  (tempera- 
ture above  which,  at  760  mm.  pressure,  the  substance  is  gaseous),  spe- 
cific gravity,  specific  heat,  and  conductivity  for  electricity.  Thus, 
sulphur  is  yellow,  has  little  odor,  crystallizes  in  the  rhombic  system, 
has  a  hardness  of  2.5  on  a  scale  of  ten,  has  the  m.-p.  115°  C.,  is  not 
perceptibly  soluble  in  water  but  dissolves  in  carbon  disulphide  (41 : 100 
at  18°),  has  the  b.-p.  448°  C.,  the  sp.  gr.  2,  the  sp.  ht.  0.18,  and  is  a  very 
poor  conductor.  In  the  first  two  chapters,  while  presenting  the  ex- 
perimental facts  required  for  our  discussion,  we  have  had  to  speak 
of  substances  very  frequently  (e.g.  pp.  26  and  32),  and  have  done  so 
always  in  the  above  sense. 

Slight  variations  from  the  standard  properties  of  the  substance  usually  indi- 
cate the  presence  of  an  impurity  homogeneously  incorporated.  The  precise  ways 
in  which  the  properties  are  affected  in  such  cases  will  be  considered  under  solu- 
tions. 

The  "  substance  "  will  be  seen  to  be  of  an  abstract  nature.  It  is  a  conception 
built  up  by  selecting  (or  abstracting)  the  properties  common  to  all  specimens. 
Hence  we  classified  chemistry  as  an  abstract-concrete  science  (p.  3).  The  bodies 
under  observation  are  concrete,  the  classification  of  the  results  is  under  conceptions 
of  an  abstract  nature  like  this  one. 

That  all  bodies  of  a  like  kind  have  many  identical  properties  is  the  most  funda- 
mental fact  in  chemistry.  Being  a  general  fact,  we  call  it  a  law,  and  word  it  as  fol- 
lows :  The  specific  properties  of  a  substance  are  constant  in  all  speci- 
mens. Experience  having  given  us  confidence  in  its  universality,  we  take  it  for 
granted  in  all  our  work  (cf.  First  characteristic,  p.  5). 

There  are  still  other  qualities  which  a  body  (or  specimen  of  mat- 
ter) may  possess.  It  has,  for  example,  a  certain  temperature,  press- 
ure, motion,  or  electric  charge.  These  we  speak  of  as  conditions  of 


36  INORGANIC   CHEMISTRY 

the  body  rather  than  properties.  In  the  use  of  the  body  they  may  be 
altered,  and  some  of  them  may  be  removed  or  added,  arbitrarily. 

Thus  there  are  three  kinds  of  qualities  to  be  considered.  The 
attributes,  like  weight,  form,  and  volume,  do  not  belong  to  substanqes 
but  to  bodies.  The  specific  properties,  like  color,  solubility,  and  odor, 
belong  by  right  to  substances,  although  we  sometimes  speak  of  them  in 
connection  with  bodies.  The  conditions,  like  temperature  and  motion, 
belong  to  neither,  for  they  can  be  altered  without  changing  either  the 
body  or  the  substance. 

Methods  of  Work  and  Observation  in  Chemistry.  —  It  is  not 

the  end  of  chemical  work  to  make  generalizations  or  laws,  like  the  char- 
acteristics of  chemical  phenomena  (pp.  5,  16,  etc.),  or  conceptions,  like 
those  dealt  with  in  the  preceding  paragraph.  These  are  simply  the 
means  by  the  help  of  which  chemical  work,  whether  it  be  investigation, 
commercial  analysis,  or  manufacturing,  may  be  carried  on  more  system- 
atically. Together  they  constitute  our  system  for  classifying  the  facts 
with  a  view  to  ready  reference.  The  sample  experiments  (pp.  10-14), 
if  reexamined,  will  show  that  we  there  employed  most  of  the  cate- 
gories of  our  classification  which  have  so  far  been  described. 

Thus,  in  the  experiment  with  iron  and  sulphur  (p.  10),*  it  was  first 
our  object  to  find  out  whether  the  bodies  had  interacted  chemically  on 
being  mixed.  To  do  this  we  endeavored  to  ascertain  whether  any  por- 
tion of  the  mixture  had  acquired  new  specific  properties  (p.  35). 
Here  we  used  the  law  (p.  7)  that  all  the  properties  of  the  products 
of  a  chemical  interaction  are  different  (pp.  5-7)  from  those  of  the  initial 
substances.  We  also  noted  the  specific  properties  of  the  substances 
concerned,  purposely  omitting  all  mention  of  quantity  and  tempera- 
ture, because  attributes  (p.  35)  like  the  former  and  conditions  (p.  35) 
like  the  latter  do  not  characterize  substances  (p.  35)  in  general,  as  dis- 
tinct from  particular  specimens,  and  cannot  be  used  for  identification. 
We  found  that  a  part  dissolved  in  carbon  disulphide  and  the  remainder 
was  all  magnetic. 

While  all  the  spr  eit;  properties,  of  which  a  few  are  mentioned  on 
p.  35,  find  application  IP  identification,  the  first  seven  in  that  list  are 

*  References  to  previous  pages  are  used  in  order  to  save  needless  repetition 
in  writing.  But  the  beginner  requires  endless  repetition  in  his  reading  and  must 
form  the  habit  of  examining,  in  conjunction  with  the  current  text,  the  parts  referred 
to.  The  passages  cited  are,  by  the  reference,  made  part  of  the  current  text,  which 
will  usually  not  be  clear  without  them.  The  same  remark  applies  to  topics  referred 
to  by  name,  and  to  be  sought  in  the  index. 


INTRODUCTORY   II  37 

most  frequently  used.  And  of  these  seven  the  solubility  in  various 
solvents  and  the  boiling-point  are  by  far  the  most  important.  This  is 
because  they  are  best  suited  for  separating  mixtures.  Thus,  here,  we 
first  removed  the  sulphur  by  dissolving  it,  and  then,  after  evaporating 
away  the  solvent,  we  completed  the  identification  by  recognizing  the 
color  and  form  of  the  residue.  Similarly  we  recognized  by  its  appear- 
ance and  magnetic  property  the  iron  that  was  left  undissolved. 

Separation  by  the  use  of  one  of  these  two  properties,  then  application 
of  the  others  for  identification,  is  the  general  order  of  procedure  in 
chemistry.  Most  of  the  other  properties  cannot  be  recognized  readily 
in  mixtures,  as  a  moment's  thought  will  show.  The  general  color  and 
the  specific  gravity  of  a  mixture,  containing  unknown  substances  in  un- 
known proportions,  for  example,  tell  us  little  about  the  corresponding 
properties  of  the  components.  Magnetic  properties  may  be  used  for 
separation,  also,  but  iron  is  almost  the  only  substance  which  shows 
them  markedly,  so  that  their  application  for  the  purpose  is  very 
limited. 

The  use  of  the  boiling-point  (temp,  of  free  vaporization)  for  sepa- 
ration was  illustrated  in  the  evaporation  of  the  carbon  disulphide  to  get 
the  sulphur  alone.  This  solvent  has  a  very  low  boiling-point  (46°  C.) 
and  therefore  a  high  vapor  pressure  at  the  temperature  of  the  room. 
By  virtue  of  this  it  evaporates  rapidly.  The  sulphur  (b.-p.  448°  C.)  is 
not  volatile  under  these  conditions  and  remains  behind. 

In  connection  with  this  investigation  we  employed  several  of  the 
common  methods  of  manipulation  used  by  the  chemist.  These  methods 
are  derived  from  the  conceptions  described  in  last  paragraph.  Thus 
we  treated  the  mixture  -with  a  solvent  (Fig.  2),  on  the  assumption 
that  if  it  was  heterogeneous  (p.  34)  the  components  would  each  behave 
as  if  alone  present.  We  then  filtered,  a  method  invented  for  dealing 
with  a  heterogeneous  mixture  consisting  of  a  solid  and  a  liquid. 
Decantation  is  often  used  in  such  cases  when  the  solid  is  specifically 
much  heavier  than  the  solvent  and  settles  readily.  We  allowed  the  car- 
bon disulphide  to  evaporate  spontaneously,  and  this  is  our  favorite 
method  of  dealing  with  a  mixture  which  is  homogeneous,  and  therefore 
would  run  through  a  filter  as  a  whole  without  suffering  separation. 
When  the  liquid  has  a  higher  boiling-point  than  50-60°  C.,  as  water  has, 
we  use  heat  from  a  steam-bath  or  Bunsen  flame  to  promote  the  evapora- 
tion. In  evaporation  we  allow  the  vapor  of  the  liquid  to  escape,  be- 
cause it  is  the  less  volatile,  dissolved  body  that  we  wish  to  examine. 
When  we  desire,  on  the  contrary,  to  examine  the  liquid,  the  vapor 


38 


INORGANIC  CHEMISTRY 


must  be  condensed.  This  method,  which  we  have  not  yet  had  occasion 
to  employ,  is  called  distillation  (Fig.  16).  The  jacket  round  the  long 
tube  is  filled  with  a  stream  of  cold  water,  which,  on  account  of  its  high 
specific  heat,  quickly  cools  and  condenses  the  vapor.  The  resulting 
liquid  is  caught  in  a  flask. 

These  methods  may  be  adapted  to  the  investigation  of  any  similar 
problem.  Thus,  gunpowder  is  made  by  the  intimate  mixing  of  sulphur, 
charcoal,  and  saltpeter  (potassium  nitrate).  If  no  chemical  interaction 
whatever  has  occurred,  a  sample  will  be  wholly  separable  into  these 
components.  If  a  partial  change  has  taken  place,  a  certain  amount  of 


FIG.  16. 

material  with  different  properties  will  be  discovered  in  the  mixture. 
If  the  change  has  been  complete,  no  portion  of  the  original  substances 
will  be  found.  We  must  first  study  the  specific  properties  of  each  of 
the  ingredients  separately,  in  order  that  a  plan  of  separation  may  be 
devised,  and  that  we  may  have  a  basis  for  comparison  with  the  products 
of  the  separation. 

The  experiments  with  mercuric  oxide  (p.  12)  and  with  silver 
nitrate  (p.  13)  simply  ring  the  changes  on  the  same  conceptions. 
When  we  heat  the  former  it  is  resolved  into  its  constituents.  We  se- 
lected the  experiment  because  the  conditions  are  such  that  separation 
of  the  products  from  each  other  and  from  the  original  material  occurs 
spontaneously  without  further  manipulation.  The  separation  depends 


INTRODUCTORY  11  39 

here  on  the  boiling-points  of  the  materials.  That  of  oxygen  is  very 
low  (  —  183°  C.)  ;  so  that,  being  a  gas  even  at  ordinary  temperatures,  it 
comes  off.  That  of  mercury  is  much  higher  (357°  C.),  but  the  heat  pro- 
duced by  the  flame  (1000°-1200°  C.  in  the  bottom  of  the  test-tube)  is 
more  than  sufficient  to  vaporize  this  product  also.  The  speedy  con- 
densation on  the  cooler  part  of  the  tube,  causing  separation  of  the  two 
products,  is  due  to  the  wide  difference  in  boiling-points.  The  part  of 
the  mercuric  oxide  which  is  still  unchanged,  on  the  other  hand,  is 
entirely  involatile  even  at  1200°  C. ;  so  that  it  does  not  mingle  with 
the  vaporized  products  at  all. 

In  the  action  of  silver  nitrate  on  sodium  chloride  (p.  13),  it  was 
solubility  that  furnished  the  means  of  separating  the  products.  The 
silver  chloride  is  practically  insoluble  in  water,  while  sodium  nitrate  is 
very  soluble.  To  recover  the  latter,  the  boiling-points  were  then  con- 
sidered, and  the  more  easily  vaporized  water  was  driven  off. 

When,  following  the  latter  experiment,  the  silver  chloride  was  exposed  to  light, 
resolution  into  silver  and  chlorine  took  place.  The  products  became  separated 
spatially  at  once,  because  chlorine  is  a  gas  (substance  of  low  b.-p.)  and  silver  is  a 
solid  (usually  a  substance  of  high  b.-p.). 

In  every  case  the  specific  properties —  color,  crystalline  form,  and  so 
forth  —  by  which  recognition  is  effected  were  mentioned. 

In  connection  with  the  foregoing,  the  following  points  should  be 
noted : 

1.  That  coincidence  in  two  or  three  specific  properties  is  generally 
sufficient  to  establish  identity. 

2.  That  we  usually  have  to  separate  the  components  of  a  mixture 
or  solution  before  identification  can  be  effected. 

3.  That  if  our  problem  were  the  making  of  a  particular  compound 
whose  constituents  were  known,  we  should  study  the  products  of  each 
attempt  by  the  above  method.    Our  object  being  to  get  the  compound  as 
a  pure  substance,  we  should  try  to  arrange  the  composition  of  each  of 
the  initial  substances  (apart  from  the  necessary  constituent)  so  that  any 
by-products  formed  should  have  .such  properties  that  we  could  readily 
effect  a  separation  of  them  from  the  desired  compound. 

4.  That  a  wide  knowledge  of  specific  physical  properties  is  required 
for  intelligent  chemical  work  (cf.  pp.  36-39  and  65). 

5.  That  our  methods  are  purposely  limited  so  as  not  to  separate 
chemically  combined,  but  only  physically  mixed,  forms  of  matter  :  after 
a  physical  individual  has  been  isolated,  and  even  then  only  if  it  has  new 


40  INORGANIC   CHEMISTRY 

properties,  and  is  not  recognized  as  a  known  substance,  we  next  proceed 
to  separate  it  into  its  chemical  constituents  so  as  to  learn  which  con- 
stituents it  contains  and  in  what  relative  proportions  by  weight. 

Summary.  —  In  this  chapter  we  have  added  considerably  to  our 
conception  of  the  scope  of  chemistry  (cf.  p.  16).  Although  our  survey 
is  by  no  means  yet  complete,  we  may  condense  our  results  as  follows : 

Chemistry  deals  with  the  changes  in  composition  and  constitution 
which  substances  undergo  and  with  the  transformations  of  energy 
which  accompany  them.  To  convert  the  isolated  facts  into  a  science  we 
classify  related  parts  under  laws,  such  as  those  of  conservation  of  mass 
(p.  17)  and  of  energy  (p.  23),  and  under  conceptions,  such  as  those  of 
chemical  energy  (p.  25),  element  (p.  31),  body  (p.  34),  substance  (pp. 
26,  32,  and  35).  We  also  distinguish  between  attributes,  specific 
properties,  and  conditions  (pp.  35-36).  In  the  last  paragraphs  we 
have  indicated  briefly  the  use  to  which  these  conceptions  and  this 
classification  are  put. 

Chemical  laboratory  work  consists  largely  in  the  separation,  recog- 
nition, and  description  of  substances.  The  importance,  especially,  of 
thorough  familiarity  with  specific  properties  and  the  influence  of  con- 
ditions (for  example,  temperature)  to  these  ends  is  shown  by  the  ex- 
amples (pp.  36-39).  The  system  of  classification  as  a  whole  is  part  of 
the  everyday  mode  of  thought  of  the  chemist,  for  thought  consists 
largely  in  comparing  and  contrasting,  and  our  system  of  classification 
furnishes  the  norm  of  this  so  far  as  chemistry  is  concerned.  Learning 
chemistry  consists,  therefore,  in  learning  this  classification  and  becom- 
ing habituated  to  its  use. 

The  influence  of  conditions  has  as  yet  been  barely  touched.  The 
attribute  of  quantity  will  form  the  basis  of  discussion  in  the  next 
chapter. 

Exercises.*  —  1.  Define  the  following  terms,  and  find  illustrations 
of  each,  other  than  those  given  on  pp.  34-39 :  mixture,  physical  com- 
ponent, chemical  constituent. 

2.  Describe  (a)  a  red-hot  rod  of  iron,  (b)  an  aqueous  solution  of 
sugar,  employing  all  the  terms  given  on  pp.  34-36  so  far  as  they  are 
applicable. 

*  The  exercises  should  in  all  cases  be  studied  with  minute  care.  They  not 
only  serve  as  tests  to  show  that  the  chapter  has  been  understood,  but  very  fre- 
quently also  call  attention  to  ideas  which  might  not  be  acquired  from  the  text  alone, 
or  assist  in  elucidating  ideas  given  in  the  text  which,  without  the  exercises,  might 
not  be  fully  grasped. 


CHAPTER  III 
INTRODUCTORY   III 

The  Law  of  Definite  Proportions,  fifth  Characteristic  of 
Chemical  Phenomena.  —  In  making  a  chemical  compound,  may  we 
use  varying  proportions  of  the  constituents  ?  The  controversy  between 
Berthollet  and  Proust,  in  which  the  former  supported  the  affirmative 
and  the  latter  the  negative  side  of  this  question,  was  one  of  the  chief 
features  of  the  chemical  history  of  the  early  part  of  the  nineteenth 
century.  The  ways  of  forming  or  decomposing  a  compound,  or  of 
carrying  out  a  more  complex  chemical  change,  may  be  varied  indefi- 
nitely. The  apparatus,  the  mode  of  experiment,  and  the  proportions 
of  the  materials,  may  be  altered  at  our  will.  But,  in  spite  of  an 
enormous  amount  of  careful  work,  no  case  of  variation  in  the  propor- 
tion of  the  constituents  actually  used  or  produced  in  a  given  chemical 
action  has  come  to  light.  If  too  much  of  one  constituent,  for  example, 
is  taken,  a  part  simply  remains  unchanged.  A  higher  temperature 
may  hasten  the  chemical  action,  but  it  does  not  affect  the  quantitative 
composition  of  the  products,  provided  the  resulting  substances  are  of 
the  same  nature.  It  was  the  work  of  Stas  (1860-65)  which  settled  the 
question  by  proving  that  even  slight  variations  cannot  be  detected.  It 
is,  therefore,  a  characteristic  of  chemical  phenomena  that :  In  every 
sample  of  each  compound  substance  formed  or  decomposed,  the  pro- 
portion by  weight  of  the  constituents  is  always  the  same.  This  state- 
ment of  fact  is  known  as  the  law  of  definite  proportions.  When  the 
composition  of  a  substance  seems  to  be  variable,  it  is  always  found  on 
closer  examination  that  mechanical  mixtures  of  some  kind  were  being 
mistaken  for  pure  substances. 

Another  form  of  statement,  which  is  a  corollary  of  this  one,  and  is  applicable 
more  directly  to  complex  chemical  actions,  is :  The  ratio  by  weight  of  any  one  of 
the  factors  or  products  of  a  chemical  change  to  any  other  is  constant. 

The  Laiv  of  Multiple  Proportions,  Sixth  Characteristic.— 

In  the  course  of  even  a  very  limited  experience  in  the  examination  of 
chemical  compounds,  we  should  be  bound  to  find  that  among  the  sub- 
stances which  we  examined  were  cases  of  two  or  more  perfectly  dis- 

41 


42  INORGANIC   CHEMISTRY 

fcinct  bodies,  made  of  the  same  elements,  but  on  analysis  found  to  con- 
tain different  proportions  of  the  constituents.  We  should  discover 
that  each  of  those  distinct  substances  had  a  constant  composition  pecul- 
iar to  itself.  Thus,  we  find  in  nature  a  beautiful  yellow  mineral 
known  as  pyrite.  A  pure  sample  of  this  contains  nothing  but  iron  and 
sulphur,  and  yet  the  composition  of  the  substance  is  entirely  different 
from  that  of  ferrous  sulphide  (p.  12).  The  percentage  of  the  constit- 
uents in  each  is  as  follows  : 

Ferrous  sulphide.       Pyrite. 
Iron  63.59  46.62 

Sulphur        36.41  63.38 

If  we  calculate  from  these  data  how  much  of  one  of  the  elements  is 
combined  with  identical  amounts  of  the  other  in  each  of  the  two  com- 
pounds, a  simple  relation  emerges  from  the  seemingly  unrelated  pro- 
portions. Thus,  taking  one  part  of  iron  in  each,  we  find  63.59 : 36.41 
:  :  l:x  (=  0.5725)  and  46.62  :  53.38  :  :  1 :  x  (=  1.145).  Thus  the 
quantities  of  sulphur  (0.5725  and  1.145)  combined  with  one  part  of 
iron,  in  ferrous  sulphide  and  pyrite  respectively,  are  in  the  ratio  1 :  2. 
The  reader  will  find  by  calculation  that  if  the  quantity  of  sulphur, 
instead  of  that  of  iron,  be  fixed  at  any  value,  the  proportions  of  iron 
in  the  compounds  will  then  stand  in  the  ratio  2  : 1. 

Again,  sulphur  burns  in  the  air,  forming  with  oxygen  a  gaseous  sub- 
stance whose  odor  is  familiar.  But  a  different  compound  of  the  two 
elements,  generally  seen  in  the  form  of  a  white  fibrous  solid,  is  much 
used  in  dye-factories.  In  the  former,  the  proportion  of  sulphur  to 
oxygen  is  almost  exactly  1 : 1  (50  per  cent  of  each),  while  in  the  lat- 
ter it  is  1 :  l£  (40  per  cent  sulphur  and  60  per  cent  oxygen).  Thus 
the  two  different  proportions  of  oxygen  combining  with  one  part  of 
sulphur  are  in  the  ratio  1 :  l£  or  2  :  3. 

The  existence  of  this  simple  relation  is  not  an  accident,  and  con- 
fined to  these  cases.  It  is  a  rule  to  which  no  exception  has  yet  been 
found.  More  complex  examples  yield  the  same  result. 

Thus  there  are  known  over  two  hundred  compounds  of  carbon  and  hydrogen, 
all  different  in  composition.  But  if  we  fix  the  quantity  of  one  of  the  elements  and 
find  the  amount  of  the  other  which,  in  each  compound,  is  combined  with  that 
quantity,  the  ratio  of  these  amounts  to  each  other  is  expressible  by  integral  num- 
bers. Take,  as  a  sample,  the  compositions  of  four  of  these  substances  :  methane, 
acetylene,  ethylene,  and  naphthalene  (moth-balls)  contain  respectively  3,  12,  6, 
and  16  parts  of  carbon  combined  with  one  part  of  hydrogen. 

These  ratios  are  not  those  of  approximately  whole  numbers.  The 
more  carefully  our  determinations  of  the  proportions  of  the  constitu- 


INTRODUCTORY  HI  43 

ents  in  the  several  compounds  have  been  made,  the  more  exactly  do 
integral  numbers  represent  the  numerical  relations  between  the  results. 
This  principle  was  discovered  by  Dalton  (1804),  and  was  embodied 
by  him  in  a  statement  known  as  the  law  of  multiple  proportions, 
which  ran  somewhat  as  follows  :  If  two  elements  unite  in  more  than 
one  proportion  forming  two  or  more  compounds,  the  quantities  of  one 
of  the  elements,  which  in  the  different  compounds  are  united  with 
identical  amounts  of  the  other,  stand  to  one  another  in  the  ratio  of  in- 
tegral numbers,  which  are  usually  small. 

The  Measurement  of  Combining  Proportions.  —  The  most 
exact  measurement  of  the  proportions  in  which  the  elements  combine 
to  form  compounds  involves  manipulations  too  elaborate  to  be  gone 
into  here.  Operations  of  the  same  nature  are  described  in  works  on 
quantitative  analysis.  One  or  two  brief  statements,  diagrammatic 
rather  than  accurate,  will  show  the  principles,  however. 

If  we  take  a  weighed  quantity  of  iron  in  a  test-tube  and  heat  it  with 
more  than  enough  sulphur  (an  excess  of  sulphur),  we  get  free  sulphur 
along  with  the  ferrous  sulphide  (p.  12),  and  no  free  iron  survives. 
We  may  remove  the  free  sulphur  by  washing  the  solid  with  carbon 
disulphide.  The  difference  between  the  weights  of  ferrous  sulphide 
and  iron  gives  the  amount  of  sulphur  combined  with  the  known 
quantity  of  the  latter. 

As  an  example  of  the  study  of  rusting,  we  may  weigh  a  small 
amount  of  copper  in  the  form  of  powder  in  a  porcelain  boat  and  pass 
oxygen  over  the  heated  metal  (Fig.  17).  The  formation  of  cupric  oxide 
takes  place  rapidly.  If  we  limit  the  oxygen,  part  of  the  copper  may 
remain  unaltered ;  if  we  use  it  freely,  the  excess  will  pass  on  unchanged. 
A  given  weight  of  copper  cannot  be  induced  to  take  up  more  than  a 
certain  amount  of  oxygen,  and  use  of  a  less  amount  simply  limits  the 
amount  of  copper  transformed  into  oxide.  The  original  weight  of 
the  copper,  and  the  increase  in  weight,  representing  oxygen,  give  us  the 
data  for  determining  the  composition  of  cupric  oxide.  The  record  of 
the  result  is  usually  made  in  the  form  of  the  quantity  of  each  con- 
stituent in  a  hundred  parts  of  the  compound,  and  is  called  the  per- 
centage composition. 

The  data  furnished  by  one  rough  lecture-experiment,  for  example, 
were  as  follows : 

Weight  of  boat  empty 3.428  g. 

Weight  of  boat  +  copper 4.278  g. 

Difference  =  weight  of  copper 0.860  g. 


44 


INORGANIC   CHEMISTRY 


Weight  after  addition  of  oxygen 4.488  g. 

Weight  without  oxygen 4.278  g. 

Difference  =  weight  of  oxygen 0.210  g. 

The  proportion  of  copper  to  oxygen,  so  far  as  this  one  measurement 
goes,  is  therefore  85  :  21. 


FIG.  17. 


To  find  the  percentage  of  each  constituent,  we  observe  that  the  pro- 
portion of  copper  is  85  :  85  4-  21,  or  fifo  of  the  whole.  That  of  the 
oxygen  is  ^V^  of  the  whole.  Thus  the  percentages  are  : 

Copper,          106  :  85::  100  :  x        x  =  80.2 
Oxygen,          106  :  21  ::  100  :  x'        x'=  19.8 

Naturally,  the  mean  of  the  results  of  a  number  of  more  carefully 
managed  experiments  will  be  nearer  the  true  proportion.  The  percen- 
tages at  present  accepted  as  most  accurate  are  79.9  and  20.1. 

In  the  case  of  mercuric  oxide,  we  may  decompose  a  known  weight 
of  the  oxide  (p.  12)  and  afterwards  weigh  the  mercury  and  ascertain 
the  oxygen  by  difference. 

Finally,  a  strip  of  the  metal  magnesium  may  be  set  on  fire  in  the  air. 
It  gives  out  a  dazzling  white  light  in  burning,  and  on  this  account  the 
powdered  metal  is  used  in  making  flash-light  powder  for  photography. 
The  product  is  magnesium  oxide,  a  white  substance,  which  partly  rises 
as  a  dense  smoke  and  partly  falls  on  the  ground.  In  a  loosely  closed 
porcelain  vessel  (Fig.  18)  the  metal  may  be  burned  slowly,  with  the 
help  of  the  heat  from  a  small  flame,  and  the  oxide  may  be  retained. 


INTRODUCTORY  III 


45 


The  weight  of  magnesium  ribbon  taken  and  the  increase  in  weight  due 
to  oxygen  give  the  data  for  calculating  the  proportions  of  the  constit- 
uents. 

The  following  figures  show  the  results  of  experiments  in  these  and 
other  simple  cases,  and  represent  the  percentage  composition  of  the 
products,  only  two  places  of  decimals  being  given  in  each  case.  They 
will  be  required  for  the  discussion  of  the  next  topic. 


(1) 


(2) 


(3) 


Cupric  oxide 
Copper,         79.9 
Oxygen,        20.1 
Mercuric  oxide 
Mercury,     92.59 
Oxygen, 
Water 
Hydrogen 
Oxygen, 


7.41 

11.18 

88.81 


(4) 


(5) 


(6) 


Cupric  sulphide 
Copper,       66.48 
Sulphur,      33.51 
Mercuric  sulphide 
Mercury,     86.18 
Sulphur,      13.81 
Hydrogen  sulphidt 
Hydrogen,     5.92 
Sulphur,      94.07 


(7)  Cupric  chloride 
Copper,        47.3 
Chlorine,     62.7 

(8)  Mercuric  chloride 
Mercury,      73.8 
Chlorine,     26.2 

(9)  Hydrogen  chloride 
Hydrogen,     2.76 
Chlorine,     97.23 


The  Law  of  Combining  Weights,  Seventh  Characteristic. — 

The  ratios  in  the  above  list  represent  the  true  proportions  by  weight  in 
the  various  compounds,  but 
naturally  the  individual  num- 
bers constituting  those  pro- 
portions have  no  chemical  sig- 
nificance whatever.  They  are 
arbitrary  values  selected  so 
that  the  constituents  of  the 
proportion  may  together  make 
100.  Each  pair  represents 
the  constant  ratio  which  is  the 
mean  result  of  numerous  ex- 
periments. 

We  begin  the  effort  to  re- 
duce these  numbers  to  order 
by  selecting  one  element  as 
our  starting-point,  and  by  tak- 
ing some  convenient  weight 
of  it  as  the  basis.  As  we 

shall  see,  it  makes  no  difference  what  choice  we  make  in  either  re- 
spect. To  avoid  waste  of  time,  we  shall,  therefore,  use  oxygen,  as  it  is 
the  element  generally  preferred  by  chemists  for  the  purpose.  The 
reason  for  this  preference  will  be  apparent  later  (see  pp.  48,  51) 


FIG.  18. 


46  INORGANIC  CHEMISTRY 

We  should  naturally  be  inclined  to  use  1  part  of  the  element  as  our 
basis.  But  our  later  steps  involve  finding  out  what  amounts  of  the 
other  elements  combine  with  this  quantity,  and  we  perceive  that  the 
amount  in  the  case  of  hydrogen  will  be  only  0.126  parts.  We  calculate 
this  from  (3) :  88.81  :  11.18  :!:«(=  0.126).  If,  however,  we  take 
8  parts  of  oxygen,  this  amount  of  hydrogen  is  also  increased  eight 
times  and  becomes  1.008.  As  no  element  is  found  to  combine  in 
smaller  proportions  than  hydrogen,  we  are  satisfied  that  a  scale  for  our 
numbers  based  on  8  parts  of  oxygen  will  not  involve  any  values  less 
than  1.  The  choice  of  scale  is  purely  one  of  convenience. 

Our  further  procedure  is  determined  by  a  desire  to  demonstrate  a 
particular  relationship,  and  so  we  must  keep  a  fixed  principle  in  mind. 
This  principle  is  that  of  calculating  a  series  of  combining  proportions 
of  which  each  succeeding  member  is  the  amount  of  some  element  which 
combines  with  the  amount  of  the  one  preceding  it  in  the  list.  Oxygen 
is  the  initial  member  of  the  series. 

From  (1)  we  calculate  the  amount  of  copper  combined  with  8  parts 
of  oxygen,  thus:  20.1 :  79.9  :  :  8  :  x  (=  31.8).  We  next  look  for  an  ele- 
ment which  combines  with  copper —  any  such,  whether  it  appeared  in 
the  above  list  or  not,  would  do — and  from  (4),  for  example,  find,  by 
the  same  method  of  calculation,  that  31.8  parts  of  copper  unite  with 
16.03  parts  of  sulphur.  Then  we  select  any  element  that  combines  in 
turn  with  sulphur,  and  from  (5)  learn  that  16.03  parts  of  sulphur  unite 
with  100  parts  of  mercury.  Next  we  observe  from  (8)  that  100  parts 
of  mercury  combine  with  35.45  parts  of  chlorine.  Finally,  from  (9)  we 
calculate  that  35.45  parts  of  chlorine  unite  with  1.008  parts  of  hydro- 
gen. Setting  these  results  down  in  order,  we  obtain  the  series: 

OXYGEN  COPPER         SULPHUR       MERCURY      CHLORINE       HYDROGEN 

8  31.8  16.03          100          35.45  1.008 

Since  each  of  these  numbers  represents  the  amount  of  the  corre- 
sponding element  which  is  actually  found  to  combine  completely  with  the 
quantities  of  two  other  elements,  we  may  call  them  combining  weights. 
It  must  be  observed,  however,  that  those  of  our  measurements  of  which 
we  have  made  use  in  deriving  them  relate  each  number  in  the  series  to 
its  two  immediate  neighbors  only. 

Now  (2)  gives  the  proportion  in  which  oxygen  combines  with  mer- 
cury. By  calculation,  7.41  :  92.59  :  :  8  :  x  (=  100),  we  find  that  8 
parts  of  oxygen  combine  with  100  parts  of  mercury.  It  must  be 
pointed  out  explicitly  that  this  result  for  mercury  and  oxygen  is  en- 


INTRODUCTORY  III 


47 


tirely  independent  of  the  previous  one  for  mercury  and  sulphur,  by 
which  the  value  100  was  originally  derived.  It  is  based  upon  the  in- 
dependent measurement  of  the  proportion  of  mercury  in  mercuric  oxide, 
and  we  could  not  possibly  have  foretold  what  this  proportion  would  be. 
It  appears  then  that  the  above  combining  weights,  at  first  measured  for 
neighboring  pairs  only,  are  also  the  weights  according  to  which  remoter 
combinations  take  place.  Further  study  confirms  this  speculation. 
Thus  (3)  shows  that  oxygen  and  hydrogen  combine  in  the  proportion 
8:  1.008,  and  (7)  gives  the  proportion  in  cupric  chloride,  and  (6)  that  in 
hydrogen  sulphide  exactly  as  they  stand  in  the  series. 

A  few  other  pairs  in  this  set  are  capable  of  uniting.  For  example, 
oxygen  and  sulphur,  as  we  have  seen,  form  two  compounds  in  which 
the  proportions  are  found  to  be  2  x  8  :  16.03  and  3x8:  16.03.  Sulphur 
and  chlorine  unite  in  the  proportions  16.03  :  2  x  35.45  and  2x16.03  : 
35.45.  Sulphur  and  copper  form  another  compound,  cuprous  sulphide, 
and  in  it  the  proportion  is  16.03  :  2  x  31.8,  instead  of  16.03  :  31.8. 
These  examples  are  also,  incidentally,  illustrations  of  the  law  of  mul- 
tiple proportions  (p.  43). 

These  relations  become  clearer  when  represented  diagrammatically. 
Thus  the  five  elements,  omitting  oxygen,  give  the  following  compounds 
and  no  others : 


-1:1- 


-1:2- 


Copper- 
(31.8) 


•1:1 Sulphur- 

(16.03) 

-2:1 


-1:1— Mercury— 1:1— Chlorine— 1:1— Hydrogen 
(106)  (35.46)  (1.008) 


-1:2- 


-2:1- 


-1:1- 

-2:1- 


It  will  be  observed  that  hydrogen  forms  stable  compounds  with  only  two  of  the 
other  four  elements  in  the  series.  If,  however,  oxygen  had  been  included,  compounds 
of  oxygen  with  all  the  other  five  would  have  demanded  recognition.  This  illustrates 
the  reason  given  below  for  the  preference  of  oxygen  as  the  fundamental  element. 

It  ought  to  be  added  that  combinations  of  three  or  more  elements  are  not  un- 
common, but  neither  does  examination  of  these  cases  reveal  anything  in  conflict 
with  the  principle  which  our  study  of  the  combining  weights  is  bringing  to  light, 
For  example,  oxygen,  chlorine,  and  hydrogen  combine  in  the  proportions  : 

2x8:  35.46  :  1.008,        6x8:  35.45  :  1.008,        and  8  x  8  :  35.45  :  1.008. 


48  INORGANIC  CHEMISTRY 

A  complete  study  of  all  known  combinations  of  all  the  elements 
shows  that  no  compound  containing  any  two  or  more  of  the  elements  is 
known,  whose  composition  cannot  be  expressed  in  terms  of  the  combining 
weights  found  in  a  series  like  the  above,  by  using  proper  integral  fac- 
tors when  necessary. 

This  statement  would  remain  true  whatever  element  we  chose  as 
the  initial  one  of  the  series  (see  Exercise  1,  p.  52),  whatever  combining 
weight  we  assigned  to  that  element,  —  for  this  affects  only  the  scale  of 
all  the  members,  not  their  relative  values,  —  and  whatever  the  order  in 
which  the  succeeding  elements  were  taken.  It  applies  to  all  the  ele- 
ments without  exception,  for  the  above  series  might  have  been  extended 
to  include  all  the  known  elements.  It  represents,  in  fact,  not  a  prop- 
erty of  our  method  of  manipulation,  but  of  chemical  combination  itself. 

It  will  now  be  seen  that,  since  the  order  is  a  matter  of  indifference, 
a  long  series  like  the  above  is  not  needed,  and  would  lead  only  to  cumu- 
lative errors.  The  exact  determinations  of  the  combining  weights  of 
most  of  the  elements  have  actually  been  made  by  direct  union  with  oxy- 
gen or  with  the  help  of  but  one  intermediate  step.  Again,  if  the 
question  had  been  one  of  mathematics,  hydrogen,  the  element  with  the 
lowest  combining  proportions,  would  have  furnished  the  basis  and  unit 
of  the  scale.  But  the  question  was  the  practical  one  of  getting  the  most 
accurate  measurements  for  the  relative  magnitudes  of  the  numbers,  so 
oxygen  was  chosen  instead.  Nevertheless,  the  value  8  was  selected  in 
order  that  the  advantage  of  having  a  mathematical  unit,  or  something 
close  to  it,  in  the  combining  weight  of  hydrogen,  might  be  retained  also. 

The  law  of  combining  weights  may  be  put  briefly  thus :  The  pro- 
portions by  weight  in  which  all  chemical  combinations  take  place  can 
be  expressed  in  terms  of  small  integral  multiples  of  fixed  numbers 
called  combining  weights,  one  for  each  element.  It  describes  what  is 
perhaps  the  most  striking  of  all  the  characteristics  of  chemical  action. 

It  is,  perhaps,  hardly  necessary  to  point  out  that  the  laws  of  definite  and  mul- 
tiple proportions  are  simply  partial  statements  whose  whole  content  is  included  in 
this  far  greater  generalization.  No  one  chemist  succeeded  in  discovering  this 
property  of  combining  weights.  The  work  of  J.  B.  Richter,  Dalton,  and  many 
others  contributed  to  it. 

Witnout  this  fact,  the  remembering  of  the  compositions  of  chemical 
substances,  necessary  as  it  is  to  the  chemist,  would  have  been  com- 
pletely beyond  the  power  of  any  ordinary  memory.  With  it,  the  task 
becomes  comparatively  simple.  It  is  only  necessary  to  decide  on  the 
best  system  of  values  for  the  combining  weights,  and  then,  regarding 


INTRODUCTORY   III  49 

the  value  of  this  for  each  element  as  the  unit  of  weight  for  that  element, 
to  express  the  proportions  of  the  element  in  every  compound  by  the 
proper  multiples.  Thus,  given  a  list  of  the  combining  weights,  one 
for  each  element,  only  the  small  integral  multiples  have  to  be  kept  in 
mind  in  connection  with  each  compound. 

The  reader  will  require  a  little  time,  however,  before  he  becomes 
accustomed  to  the  use,  not  of  a  single  unit  of  weight,  but  of  a  different 
one  for  each  element.  Chemistry  is  the  only  science  in  which  the 
physical  unit  of  weight,  which  is  the  same  for  all  materials,  is  not 
employed  for  every  purpose.  The  physical  manipulations  of  the 
chemist  are  carried  out  with  the  use  of  physical  units,  but  the  chemi- 
cal results  are  expressed  in  terms  of  individual  unit  quantities  of  the 
several  elements,  the  combining  weights. 

The  individual  units  actually  used  for  each  element  are  not  in  all 
cases  identical  with  those  we  have  given.  The  final  values  will  be 
discussed  in  the  next  section  but  one. 

Most  of  the  first  exact  determinations  of  combining  proportions  were  made  by 
Berzelius  before  1830.  It  should  be  added  that,  while  the  combining  weights,  with 
the  exception  of  that  of  oxygen  which  is  the  standard,  are  never  actually  whole 
numbers,  although  they  often  approach  such  integral  values,  the  integers  used  to 
multiply  them,  when  they  are  employed  to  express  combining  proportions,  are  so 
most  exactly.  Even  in  determinations  by  methods  of  the  highest  refinement,  the 
factors  to  be  used  in  multiplying  the  combining  weights  are  always  found  to  di- 
verge from  whole  numbers  by  amounts  within  the  known  errors  of  the  method  of 
measurement. 

Equivalents.  —  The  combining  weights  may  be  viewed  from  an- 
other standpoint.  The  relation  between  the  first  three  members  of 
the  series  discussed  in  the  last  section,  for  example,  may  be  stated 
thus  :  16.03  parts  of  sulphur  and  8  parts  of  oxygen  combine  with  iden- 
tical amounts  of  copper,  namely,  31.8  parts.  Either  of  them  will  satisfy 
the  same  amount  of  this  third  element.  In  fact,  they  are  equivalent 
for  the  purpose  of  combination.  Similarly,  31.8  parts  of  copper  and 
100  parts  of  mercury  are  equivalent,  because  either  will  hold  16.03 
parts  of  sulphur  in  combination. 

Now,  copper  will  actually  displace  mercury  from  combination  with 
a  third  element.  The  copper  will  combine  with  the  third  element,  and 
the  mercury  will  be  set  free.  So,  if  we  know  the  combining  weight  of 
copper,  but  not  that  of  mercury,  we  can  make  a  measurement  of  the 
amount  of  mercury  displaced  by  31.8  parts  of  copper.  This  amount 
(100  parts),  which  we  the  a  call  the  "  equivalent,"  will  be  also  the  re- 


50 


INORGANIC   CHEMISTRY 


quired  combining  weight.  It  will  be  so  because  this  amount  of  mercury 
was  formerly  in  combination  with  that  quantity  of  the  third  element 
with  which  31.8  parts  of  copper  are  now  united.  The  terms  combin- 
ing weight  and  equivalent  both  refer  to  the  same  members,  and  the 
latter  simply  emphasizes  a  particular  experimental  method.  Follow- 
ing the  custom  of  chemists,  we  shall  hereafter  refer  to  the  combining 
weight  on  the  scale  oxygen  =  8  as  the  equivalent  of  the  element.  The 
equivalent  weight  of  an  element  is  the  quantity  of  the  element  which 
combines  with  or  displaces  8  parts  of  oxygen  or  1.008  parts  of  hydrogen. 

Atomic  Weights.  —  The  chemist  frequently  uses  the  idea  of 
equivalents  and  the  values  we  have  given  them.  But  far  more  often 
he  employs  a  slightly  differing  set  of  numbers,  which,  for  reasons  that 
will  appear  in  a  subsequent  chapter,  he  calls  atomic  weights.  The  fol- 
lowing list  shows  the  elements  whose  equivalents  we  have  been  dis- 
cussing, along  with  one  or  two  others,  added  by  way  of  furnishing  a 
fair  sample,  and  gives  both  sets  of  weights  for  the  purpose  of  com- 
parison : 


ELEMENT. 

EQUIVA- 
LENTS (COM- 
BINING 

WEIGHTS). 

ATOMIC 
WEIGHTS. 

ELEMENT. 

EQUIVA- 
LENTS (COM- 
BINING 

WEIGHTS). 

ATOMIC 
WEIGHTS. 

Oxygen  . 

8 

16 

Iron  .     .     . 

27.95 

55.9 

Copper  . 

31.8 

63.6 

Magnesium 

12.18 

24.36 

Sulphur 

16.03 

32.06 

Carbon  .     . 

3.00 

12.00 

Mercury 

100.0 

200.0 

Aluminium 

9.03 

27.1 

Chlorine 

35.45 

35.45 

Sodium  .     . 

23.05 

23.05 

Hydrogen 

1.008 

1.008 

Bromine     . 

79.96 

79.96 

It  will  be  seen  that  some  equivalents  have  been  multiplied  by  two, 
the  first  four  and  those  of  iron  and  magnesium,  for  example ;  some 
have  been  multiplied  by  three,  like  that  of  aluminium ;  some  by  four, 
like  that  of  carbon ;  and  some  remain  unchanged,  like  those  of  chlo- 
rine, hydrogen,  sodium,  and  bromine. 

The  reasons  for  this  manipulation  of  the  simple  equivalents  found 
by  experiment  is  based  upon  theoretical  considerations.  As  it  is  im- 
possible, until  we  reach  certain  important  facts  which  cannot  be  intro- 
duced here,  to  explain  these  considerations,  the  discussion  of  the  rea- 
sons for  the  changing  of  the  numbers  will  be  postponed  until  after 
these  facts  are  before  us.  Suffice  it  to  say  that  great  advantages  are 
found  to  attach  to  these  modifications  in  the  values.  The  step  from 
equivalents  to  atomic  weights  (q.v.~)  is  taken  before  the  justification  of 


INTRODUCTORY  III  61 

it  can  be  given,  because  otherwise  formulae  (see  next  chapter)  could 
not  be  used  in  the  earlier  chapters,  and  so  the  advantages  their  em- 
ployment offers  would  be  sacrificed. 

A  little  thought  will  show  that  the  atomic  weights  have  all  the 
properties  which  we  have  shown  to  belong  to  the  combining  weights 
(equivalents).  The  atomic  weight  is  the  unit  of  weight  (p.  49)  actu- 
ally used  in  expressing  the  proportions  of  each  element  in  all  its 
compounds.  The  integral  factors  are,  of  course,  different  from  those 
which  would  be  employed  in  expressing  the  composition  of  the  same 
substance  in  terms  of  equivalents,  because  many  of  the  latter  have 
been  multiplied  by  small  integers  already  in  course  of  being  made  into 
atomic  weights.  But  the  multiplication  has  in  every  case  been  by  an 
integer,  so  that  no  change  in  the  properties  of  the  numbers  has  oc- 
curred. 

To  the  reasons  given  above  for  the  choice  of  oxygen  as  the  funda- 
mental element,  and  the  value  8  for  its  combining  or  equivalent  weight, 
one  other  may  now  be  added.  The  majority  of  the  atomic  weights, 
calculated  on  this  basis  from  the  experimental  results,  fall  so  close  to 
being  integers  that  the  nearest  round  numbers  are  exact  enough  for 
ordinary  use.  Thus  in  the  above  list  nine  of  the  twelve  atomic  weights 
are  within  0.1  of  the  nearest  whole  number.  This  convenience  disap- 
pears when,  for  example,  hydrogen  with  the  value  1  is  made  the  basis. 

The  reader  will  inevitably  find  difficulty  at  first  in  thoroughly 
grasping  the  significance  of  these  numbers.  It  may,  therefore,  be  of 
some  assistance  if  a  hint  is  thrown  out  which  will  suggest  a  concrete 
basis  for  this  curious  property.  These  numbers  appear  to  mean  that, 
when  we  wish  to  make  a  chemical  compound,  we  may  choose  any  two 
elements  from  the  list,  and,  if  it  is  found  that  they  can  combine  at  all, 
we  have  only  to  take  the  atomic  weights,  worked  out  from  other  com- 
binations of  each  element,  and  we  shall  find  that  they  will  exactly 
suffice  for  this  case  of  chemical  union.  If  complete  combination  of 
both  materials  does  not  take  place,  then  trial  will  quickly  show  what 
multiples  of  the  atomic  weights  will  result  in  this.  The  situation 
seems  to  suggest  that  the  constructing  of  chemical  compounds  depends 
upon  the  putting  together  of  ready-made  "parts,"  like  those  of  a 
watch  or  a  bicycle.  The  parts  seem  to  be  "  interchangeable,"  and 
each  element  seems  to  be  furnished  to  us  by  nature  in  ready-made 
packets  suitable  for  application  in  building  up  any  chemical  structure. 

A  complete  list  of  atomic  weights  is  printed  on  the  inside  of  the 
cover  at  the  back  of  this  book. 


52  INORGANIC   CHEMISTRY 

Summary.  —  This  chapter  adds  an  important  item  to  our  state- 
ment of  the  scope  of  the  science  (cf.  p.  40),  which,  therefore,  now 
reads  as  follows :  Chemistry  deals  with  the  quantitative  study  of  the 
changes  in  composition  and  constitution  which  substances  undergo  and 
with  the  transformations  of  energy  which  accompany  them.  To  ex- 
press the  quantitative  relations  which  are  observed,  a  different  unit  of 
weight  is  employed  for  each  element,  and  is  known  as  the  atomic 
weight  of  the  element. 

There  are  other  important  characteristics  of  chemical  phenomena 
mostly  concerned  with  the  conditions  (p.  35),  but  those  which  have 
been  given  are  sufficient,  for  the  present,  to  guide  us  in  the  systematic 
study  of  the  behavior  of  the  elements  and  their  chief  compounds. 

It  may  not  be  out  of  place  to  indicate  which  are  the  most  important  conditions. 

The  first  condition  whose  influence  we  are  likely  to  notice  in  chemical  work  is 
that  of  temperature.  The  accelerating  effect  of  rise  in  temperature  on  the  speed  of 
all  chemical  changes  (see  Chap,  v),  and  van  't  Hoff's  law  (q.v.)  in  regard  to  the 
effect  of  temperature  on  the  direction  of  chemical  change,  describe  the  most  impor- 
tant characteristics  of  this  influence. 

The  second  condition  whose  effects  we  continually  observe  is  that  of  concentra- 
tion. This,  and  not  chemical  affinity,  as  many  suppose,  determines  chemical  be- 
havior in  the  majority  of  familiar  actions.  It  is  described  by  the  law  of  concentra- 
tion (o;.i>.),  or  "  mass  action,"  as  it  is  often  inappropriately  called  to  the  great 
detriment  of  clearness.  Erin's  method  of  obtaining  oxygen  furnishes  the  first  con- 
spicuous case  of  the  influence  of  this  condition  which  we  shall  encounter.  If  this 
and  many  other  examples  are  passed  over  without  discussion,  it  is  only  because 
we  must  wait  until  much  chemical  experience  has  been  gained  before  this  principle 
can  be  understood.  Pressure  is  the  special  name  for  concentration  in  gases. 

A  third  condition  of  great  importance  in  many  —  perhaps  most  —  chemical 
actions  is  the  presence  of  a  catalytic  agent  (q.v.). 

Exercises.  —  1.  Starting  with  mercury  as  the  basal  element  and 
100  as  its  combining  weight,  calculate  from  the  data  on  p.  45  a  series 
of  combining  weights  for  the  other  five  elements.  Then  show  that  this 
series  has  the  same  properties  as  that  discussed  in  the  text. 

2.  There  is  another  oxide  of  copper,  namely,  cuprous  oxide,  which 
differs   in  properties  and  composition  from  cupric  oxide.     In  it  the 
ratio  of  copper  to  oxygen  is  2  x  31.8 :  8.     If  this  compound  had  been 
used  in  deriving  the  combining  weight  of  copper  on  p.  46,  what  effect 
would  this  procedure  have  had  on  the  rest  of  the  series  ?     Would  the 
properties  of  the  series  have  been  affected  ? 

3.  Calculate  the  equivalent  (combining)  weight  of  iron  from  the 
composition  of  ferrous  sulphide  (p.  42). 


CHAPTER  IV 
INTRODUCTORY  IV 

A  CONSIDERATION  of  the  contents  of  the  foregoing  chapter  will  show 
that  the  complete  description  of  a  chemical  change  must  be  exceed- 
ingly involved.  In  a  moderately  complex  action,  such  as  that  of 
sodium  chloride  upon  silver  nitrate,  we  should  say  that  sodium  chlo- 
ride, composed  of  one  atomic  weight  each  of  sodium  and  chlorine,  when 
brought  in  contact  with  silver  nitrate,  composed  of  one  atomic  weight 
each  of  silver  and  nitrogen  and  three  atomic  weights  of  oxygen,  gave 
silver  chloride,  composed  of  one  atomic  weight  each  of  silver  and 
chlorine,  and  sodium  nitrate,  composed  of  one  atomic  weight  each  of 
sodium  and  nitrogen  and  three  atomic  weights  of  oxygen.  Such  a 
statement,  while  it  would  give  all  the  facts  from  the  quantitative 
point  of  view,  would  be  difficult  to  grasp  and  lacking  in  perspicuity. 

Symbols,  Formulae,  and  Equations.  —  In  order  to  represent  the 
nature  of  a  chemical  change  in  a  form  which  may  be  taken  in  at  a 
glance,  the  chemist  is  in  the  habit  of  using  certain  symbols,  first  intro- 
duced by  Berzelius.  Thus,  the  letters  Ag  represent  one  atomic  weight 
(i.e.,  107.93  parts)  of  silver  (argentum),  and  O  represents  one  atomic 
weight  (i.e.,  16  parts)  of  oxygen.  In  other  words,  the  symbol  of  an 
element  means  one  chemical  unit  weight  of  the  element.  Since  many 
elements  begin  with  the  same  initial,  two  letters  have  frequently  to  be 
used  to  distinguish  them.  When  the  names  of  the  elements  are  not 
the  same  in  all  languages,  resort  is  frequently  had  to  Latin.  Thus, 
Cu  stands  for  one  atomic  weight  of  copper  (cuprum),  Fe  is  used  for 
iron  (ferrum),  Hg  for  mercury  (hydrargyrum).  From  German  we  have 
Na  for  sodium  (natrium)  and  K  for  potassium  (kalium).  To  repre- 
sent a  compound,  the  symbols  of  the  elements  which  it  contains  are 
placed  side  by  side,  small  numbers  indicating  multiples  of  the  atomic 
weights  where  they  occur.  Thus,  sodium  chloride  is  represented  by 
the  symbols  KaCl,  silver  nitrate  by  the  symbols  AgN08.  A  combina- 
tion of  symbols  is  called  a  formula.  The  symbols  of  a  formula,  taken 
by  themselves,  do  not  stand  for  any  definite  quantity  ;  each  is  one 

63 


54  INORGANIC   CHEMISTRY 

factor  of  a  proportion.  Ag  means  the  proportion  of  107.93  parts  of 
silver  to  the  proportions  of  the  other  elements  represented  by  the  other 
symbols  which  may  be  connected  with  it. 

The  value  of  these  symbols  lies  not  only  in  the  compact  way  in 
which  the  resulting  formulae  present  the  composition  of  compounds, 
but  also  in  the  use  which  may  be  made  of  them  in  showing  at  a  glance 
the  details  of  a  chemical  change.  The  chemical  action  just  mentioned 
appears  as  follows  : 

NaCl  +  AgN08  -»  AgCl 


This  expression  contains  all  that  was  conveyed  by  the  words  which 
were  written  out  in  full  above.  The  arrow  indicates  that  the  materials 
on  the  left-hand  side  pass,  in  the  chemical  transformation,  into  those 
on  the  right-hand  side.  Such  symbolic  expressions  are  called 
equations. 

One  other  variation  is  in  frequent  use.  If  we  could  make  pyrite 
(p.  42)  by  the  union  of  iron  and  sulphur,  we  should  require  two 
atomic  weights  of  the  latter  to  one  of  the  former.  The  equation 
would  run  thus  : 

Fe  +  2S  ->  FeS2. 

It  will  be  observed  that  we  employ  the  form  2S  before  combination 
and  S2  (in  FeS2)  after  it.  The  reasons  for  this  usage  will  become  clear 
as  we  proceed.  We  note  simply  that  2S  means  2  separate  atomic 
weights  of  sulphur,  as  3FeS2  would  mean  three  separate  formula- 
weights  of  pyrite.  The  same  substance  (cf.  p.  32),  sulphur,  might 
appear  as  5S  or  8S  in  other  equations,  according  to  the  proportion 
needed.  But  FeS2  is  a  group  of  three  atomic  weights  united  chemi- 
cally. The  substance  pyrite  never  contains  anything  but  two  atomic 
weights  of  the  element  sulphur,  and  its  formula  is  invariable.  Thus 
the  regular  integers  multiplying  the  atomic  weights  in  the  composi- 
tion of  a  particular  compound  are  written  after  the  symbols  of  the 
elements,  while  more  arbitrary  factors  which  change  from  one  use  of 
the  substance  to  another  are  written  in  front. 

We  shall  find  later  that  there  are  two  substances  containing  nothing  but 
oxygen,  and  that  each  is  a  compound  of  the  element  with  itself.  The  molecular 
formulae  of  these  two  are  O2  (oxygen)  and  O3  (ozone).  Thus  0  or  2O  or  3O 
would  all  be  used  for  'different  proportions  of  the  substance  O,  if  such  a  substance 
were  known,  and  O  would  be  used  for  a  substance  made  of  oxygen,  but  different 
from  oxygen  or  ozone.  Molecular  formulae  will  not  be  employed  here  until  after 
Avogadro's  hypothesis  has  been  discussed.  They  will  then  be  used  exclusively. 

The  object  in  writing  a  series  of  formulae  in  the  above  manner  is  to  show 


INTRODUCTORY  IV  55 

that  the  system  upon  the  left-hand  side,  consisting  of  certain  substances  whose 
composition  and  properties  we  know,  is,  under  the  conditions  of  the  experiment, 
unstable,  and  changes  into  the  system  upon  the  right-hand  side,  whose  nature 
we  also  know.  The  materials  on  the  two  sides  are  essentially  different,  for  the 
transformation  represented  is  a  chemical  change.  It  is  somewhat  anomalous, 
therefore,  that,  to  connect  two  sets  of  things  which  are  essentially  different,  the 
sign  =  is  usually  employed.  To  call  this  a  chemical  equation  is  still  more  anoma- 
lous, since  it  is  precisely  in  the  chemical  point  of  view  that  the  difference  between 
the  two  sides  is  most  strongly  to  be  emphasized.  It  represents  two  sets  of  things 
which  are  different,  and  not  alike  chemically.  The  physical  properties  of  the  two 
sets  of  substances  are  likewise  totally  unlike.  There  is  only  one  respect  in  which 
the  materials  on  the  two  sides  agree,  and  that  is  that  their  mass  is  not  different. 
This  is,  however,  merely  an  example  of  the  law  of  conservation  of  matter,  and 
need  not,  therefore,  be  specially  commemorated  in  the  form  in  which  we  write 
every  equation.  It  may  be  assumed  that  the  equality  in  mass  holds  for  all  chemi- 
cal changes  until  some  case  where  it  does  not  hold  shall  have  been  discovered. 
Above  all  it  must  be  remembered  that  the  chemical  equation  is  not  an  algebraic 
expression;  it  is  not  subject  to  the  rules  of  algebra.  It  is  a  brief  expression,  in 
terms  of  the  atomic  weights,  of  the  distribution  in  kind  and  quantity  of  the  con- 
stituents of  a  system  before  and  after  chemical  change. 


Making  Formulae.  —  To  make  the  formula  of  a  compound  sub- 
stance, assuming  the  formula  to  be  unknown,  two  kinds  of  informa- 
tion are  required.  We  ascertain  (1)  by  measurement  the  proportion 
by  weight  of  the  constituents  in  the  compound.  We  require  also 
(2)  to  know  the  chemical  unit  weights  —  the  atomic  weights  —  which 
have  been  accepted  by  chemists  for  each  constituent  element.  •  By 
factoring  the  terms  of  the  first  proportion  so  that  one  factor  in  each 
case  is  the  atomic  weight,  we  discover  whether  multiples  of  the  atomic 
weights  will  be  required  to  represent  the  composition  of  the  sub- 
stance, and  if  so  what  these  must  be.  An  illustration  will  make  the 
process  clear. 

Suppose  the  problem  is  to  make  the  formula  of  dried  rust.  By 
weighing  before  and  after  the  change,  we  get  the  weight  of  the  iron 
and  of  the  corresponding  amount  of  oxygen  in  the  rust  it  produces. 
If  we  took  2  g.  of  iron  we  should  get  about  2.86  g.  of  rust.  So  that 

o 

the  proportion  of  iron  to  oxygen  is    —  -  •        Now,  in  the  formula, 

0.86 

the  same  ratio  must  be  represented  by  means  of  multiples  of  the 
atomic  weights  (p.  51).  We  therefore  divide  the  quantity  of  each 
element  by  the  corresponding  atomic  weight.  This  gives  us  the 
factors  by  which  the  atomic  weights  are  to  be  multiplied.  The 
atomic  weights  are  55.9  and  16  respectively :  2  -r-  55.9  =  0.0358, 


56  INORGANIC   CHEMISTRY 

and    0.86  -j-  16  =  0.0537.       The     proportion    — ?-      then    becomes 

55.9  x  0.0358  ,,. 

•  Now  this  proportion  must  be  capable  of  expres- 
sion in  terms  of  integral  multiples  of  the  atomic  weights.  We 
find  that  the  greatest  common  measure  of  the  two  factors  is 
0.0179.  Dividing  above  and  below  by  this,  we  obtain  the  ratio 

— - — — —  •  Substituting  the  symbols  for  the  atomic  weights,  the 
16.0  X  3 

proportion  appears  as  -^-  ^—  ,  and  the  formula  is   therefore   Fe2O3. 

It  is  obvious  that  setting  the  symbols  down  side  by  side  is  not 
sufficient.  We  must  determine  by  measurement  the  factors  by 
which  they  are  to  be  multiplied. 

Making  Equations. — To  make  the  equation  representing  a 
chemical  change  we  note,  (1)  what  substances  were  used,  and  ascer- 
tain, by  study  of  their  properties,  what  substances  were  formed. 
Then  we  learn  (2)  the  formulae  of  the  substances  used  and  pro- 
duced. This  we  do  either  by  measurement  and  calculation,  as 
shown  above,  or  we  find  in  the  text-book  the  formulas  as  they  have 
been  determined  by  the  experimental  work  of  chemists.  From  these 
we  prepare  (3)  a  skeleton  equation  which,  in  the  instance  discussed, 
would  appear  thus:  Fe  -f  0  — >  Fe203.  We  are  careful  to  place  the 
initial  substances  on  the  left,  and  to  point  the  arrow  towards  those 
which  are  produced.  Finally  (4)  we  "  balance  the  equation  "  by  plac- 
ing the  proper  coefficients  before  the  formulae.  This  last  operation 
requires  experience  for  its  rapid  performance.  A  good  rule  is  to 
begin  by  picking  out  that  one  of  the  formulae  which  contains  the 
largest  number  of  atomic  weights,  no  matter  upon  which  side  it 
appears.  Here,  this  formula  is  Fe203.  We  then  reason  that,  to 
obtain  Fe2  we  require  2Fe,  and  to  obtain  O3  we  require  30,  and 
accordingly  we  place  these  coefficients  before  the  appropriate  symbols, 
thus: 

2Fe  +  30  ->  Fe2O3. 

It  is  hardly  necessary  to  add  that  a  chemical  equation  gives  the 
proportions  of  the  materials  and  nothing  more.  The  physical  condi- 
tions, for  example,  whether  the  substances  are  dissolved  in  a  liquid,  or 
are  in  the  state  of  gas,  or  are  at  a  high  temperature,  have  no  place  in 


INTRODUCTORY    IV  57 

it.  The  physical  properties  of  the  substances  concerned,  and  also 
the  energy  in  the  form  of  heat  or  electricity  which  may  appear  or  dis- 
appear in  the  process,  are  likewise  left  entirely  out.  A  question  in 
regard  to  the  nature  of  a  particular  chemical  change  demands  in 
answer  a  full  statement  of  all  these  things.  The  equation  is  therefore 
an  essential  part,  but  only  a  part,  of  such  a  statement. 

That  the  formulas  and  equations  can  deal  only  with  the  material  part  of  the 
substances  undergoing  change,  and  not  with  their  energy  (p.  26),  is  shown  by  a 
moment's  consideration.  Consistency  is  to  be  secured  only  by  holding  that  the 
symbol  S,  whether  alone  or  combined  with  others,  stands  for  the  matter-part  of  the 
sulphur  (for  the  element,  in  fact,  see  p.  32).  It  is  32  parts  by  weight  of  sulphur- 
matter.  Only  in  this  way  does  it  preserve  the  same  significance  on  both  sides  of 
the  equation  S  +  O2  — ->  SO2.  If  S  on  the  left  side  stood  for  the  free  substance  sul- 
phur, then  it  would  stand  for  32  parts  of  sulphur-matter  plus  the  appropriate 
amount  of  energy.  In  this  case  the  S  on  the  right  side  would  have  a  different  sig- 
nification, and  represent  a  less  amount  of  energy.  This  is  only  the  beginning  of 
the  difficulty,  for  we  then  find  that  S  in  H2SO4  represents  the  same  weight  of 
sulphur-matter  with  still  another  proportion  of  energy,  and  S  has  as  many  inter- 
pretations as  there  are  formulae  in  which  it  occurs.  Clearly,  all  references  to« 
energy  should  be  rigidly  excluded  from  equations,  and  thermochemical  data  can 
never  be  given  in  connection  with  them  without  complete  sacrifice  of  consistency. 
In  this  book,  however,  the  habit  of  writing  thermochemical  equations,  being 
universal,  is  frequently  followed  when  thermochemical  data  are  given. 

Units  of  Measurement  in  Chemical  Work.  —  In  chemical 
work  temperatures  are  invariably  measured  on  the  Centigrade  scale. 
The  temperature  of  a  mixture  of  ice  and  water  is  the  zero  point.  The 
temperature  of  the  steam  which  rises  from  water  boiling  under  a 
pressure  of  one  atmosphere  is  represented  by  100°.  The  interval 
between  those  two  points  is  divided  into  one  hundred  equal  parts. 

For  the  expression  of  length,  weight,  and  volume,  the  metric 
system  is  employed.  The  unit  of  this  system  is  the  meter,  which  is 
subdivided  into  decimeters,  centimeters  (cm.),  and  millimeters  (mm.). 
For  small  measurements  the  last  subdivision  is  taken  as  the  unit.  A 
cubic  centimeter  (c.c.)  is  the  unit  of  volume  for  small  measurements. 
For  larger  ones  the  liter,  which  contains  1000  cubic  centimeters,  is 
used.  The  unit  of  weight  is  that  of  one  cubic  centimeter  of  water  at 
4°,  the  temperature  of  maximum  density.  This  is  called  the  gram.* 

*  In  point  of  fact,  the  gram  is  the  one-thousandth  part  of  the  weight  of  the 
standard  kilogram  kept  in  Paris.  This  differs  from  the  weight  of  1  c.c.  of  water 
at  4°  by  less  than  0.01  per  cent. 


58  INORGANIC  CHEMISTRY 

For  larger  amounts  of  material  the  kilogram,  which  contains  1000 
grams  (1000  g.),  is  frequently  employed.  The  meter  is  equal  to  about 
39^  inches  in  ordinary  measures,  and  the  centimeter  is  very  nearly  f  of 
an  inch.  One  liter  is  about  -fa  of  a  cubic  foot  and  contains  61  cubic 
inches  or  35  fluid  ounces.  One  hundred  grams  is  about  3J  ounces 
avoirdupois,  and  one  ounce  equals  28.35  grams. 

Calculations  in  Chemistry,  —  In  the  laboratory  it  is  frequently 
desired  that  we  should  know  what  amount  of  some  substance  may  be 
obtained  by  a  given  chemical  action  from  another,  or  what  amount  of 
material  must  be  used  to  obtain  the  desired  amount  of  some  product. 
This  information  is  readily  accessible,  since  measurements  of  quantity 
in  connection  with  most  chemical^changes  are  on  record.  The  simplest 
and  most  easily  handled  form  of  this  record  is  found  in  the  formulae 
of  compounds,  and  in  the  equations  representing  the  changes  which 
they  undergo.  It  is  most  convenient,  therefore,  when  a  question  of 
this  kind  occurs,  to  ascertain  and  write  doivn,  first,  the  equation. 
Having  then  before  us  the  information  in  regard  to  the  quantities  in 
the  most  condensed  form,  we  may  use  such  parts  of  this  information 
as  are  required  for  the  problem  in  hand. 

Suppose,  for  example,  that  the  question  is  in  regard  to  the  weight 
of  oxygen  which  may  be  obtained  from  120  g.  of  mercuric  oxide.  We, 
first,  write  down  the  equation.  Next,  if  the  numbers  are  not  familiar 
to  us,  we  ascertain  the  atomic  weights  (see  Table).  These  we  then 
place  below  the  symbols  by  which  they  are  represented,  thus  : 

HgO  -»  Hg  +  0 
200+16        200        16 

Finally,  we  read  the  equation  in  its  expanded  form,  and  bring  it  into 
relation  with  the  problem  in  hand,  thus  :  One  formula-weight,  or  216 
parts  by  weight,  of  mercuric  oxide  give  one  atomic  weight,  16  parts, 
of  oxygen  ;  and  the  question  is  :  What  weight  of  oxygen  will  be 
obtained  from  120  grams  of  the  oxide  ?  The  answer  may  be  stated  by 

120 
simple  proportion  :    216  :  16  :  :  120  :  a,  or  16  X  «  ^  =  Answer.     The 


reader  must  conquer  a  tendency  to  speak  of  the  symbol  0  as  represent- 
ing "  1  part  "  of  oxygen  :  it  stands  for  16  parts.  The  word  u  part  " 
refers,  not  to  chemical  units,  but  to  physical  units  exclusively. 


INTRODUCTORY  IV  59 

Take  now  a  less  simple  case.  What  weight  of  silver  remains 
when  5  g.  of  silver  nitrate  is  heated  strongly?  The  products  are 
silver,  nitrogen  tetroxide,  and  oxygen.  We  first  write  the  equation, 
and  next  append  to  it  the  weights  it  represents : 

AgN03        ->         Ag       +       N02      +      O 
107.93  +  14  +  3x16  107.93  14  +  2x16  16 

169.93 

Then  we  read  the  equation  in  connection  with  the  problem :  If 
169.93  g.  of  silver  nitrate  yield  107.93  g.  of  silver,  how  much  of  the 
latter  will  be  furnished  by  5  g.  of  the  former  ?  169.93  :  5  :  :  107.93  :  x. 
Observe  that  O3  stands  for  3  x  16 :  the  atomic  weight  is  multiplied 
by  the  subscript  numbers,  where  such  numbers  occur. 

One  other  point  demands  illustration.  What  weight  of  chlorine  is 
required  to  combine  with  10  g.  of  mercury  and  furnish  mercuric 
chloride  ? 

Hg         +         2C1        -+         HgCl2 
200  2  x  36.45  200  +  2  x  36.45 

If  70.9  g.  of  chlorine  combine  with  200  g.  of  mercury  to  form  mer- 
curic chloride,  what  weight  of  chlorine  will  combine  with  10  g.  of 
mercury  ?  70.9  :  200  : :  x  :  10.  Observe  that  the  coefficient  in  2C1 
multiplies  the  symbol,  and  therefore  the  weight  for  which  the  symbol 
stands.  Whenever  a  coefficient  precedes  a  formula  (as  in  3AgN03), 
it  multiplies  the  whole  formula  and  therefore  the  entire  formula- 
weight. 

It  will  be  noticed  that  not  all  the  data  which  we  have  written 
down  are  necessarily  used.  In  general,  only  two  of  the  three  or  more 
weights  which  the  equation  represents  will  be  required. 

Beware  of  the  two  errors  which  lead  most  frequently  to  miscalcu- 
lation :  Do  not  omit  to  write  out  the  equation,  and  do  not  omit  to  apply 
the  expanded  equation  to  the  problem  in  hand. 

Exercises.  —  1.  What  weight  of  mercury  is  obtained  from  120  g. 
of  mercuric  oxide  ? 

2.  What  weight  of  mercuric  oxide  will  furnish  20  g.  of  oxygen  ? 

3.  What  weight  of  silver  chloride  is  obtained  from  50  g.  of  silver 
nitrate  (p.  54)  ? 

4.  What  weight  of  rust  may  be  obtained  from  10  g.  of  oxygen  ? 

5.  How  much  silver  is  contained  in  100  g.  of  an  impure  specimen 
of  silver  chloride  which  is  33  per  cent  sand  ? 


60  INORGANIC   CHEMISTRY 

6.  If  26  g.  of  mercurous  oxide  are  required  to  give,  by  heating, 
1  g.  of  oxygen,  what  is  the  formula  of  the  substance  ? 

7.  What  are  the  formulae  of  the  substances  possessing  the  follow- 
ing percentage  composition  ?     The  percentages  are  to  be  divided  by 
the  atomic  weights  just  as  were  the  actual  weights  in  the  illustration 

on  p.  55. 

I  II  III 

Magnesium,     25.57     Sodium,     32.43    Potassium,    26.585 

Chlorine,          74.43    Sulphur,    22.55    Chromium,    35.390 

Oxygen,    45.02    Oxygen,         38.025 

8.   What  are  the  percentage  compositions  of  substances  possessing 
the  following  formula :  Mn3O4,  KBr,  FeS04  ? 


CHAPTER  V 


OXYGEN 

Historical  and  Introductory.  —  Almost  one-quarter  of  the  at- 
mosphere, by  weight,  is  free  oxygen.  Water  contains  nearly  89  per  cent 
of  oxygen  in  combination,  and  this  element  constitutes  about  50  per 
cent  of  common  materials  like  sandstone,  limestone,  brick,  and  mortar. 
On  account  of  its  predominance  over  other  elements  in  quantity,  and 
the  exceptional  capacity  which  it  exhibits  for  forming  compounds  with 
a  great  variety  of  other  elements,  the  systematic  study  of  chemistry 
may  conveniently  be  begun  with  oxygen. 

While  many  elements  which  are  less  easily  obtainable  than  oxygen 
have  been  recognized  as  distinct  substances  for  many  centuries,  oxygen 
did  not  attain  this  position  until  it  was 
first  prepared  by  Priestley  in  1774. 
The  reason  of  this  was  that  gases  are 
not  so  easy  to  handle  and  distinguish  as 
solids  or  liquids,  and  consequently  very 
slow  progress  was  made  in  the  study  of 
them.  Priestley  was  particularly  in- 
terested in  examining  the  nature  of  the 
gases  which  were  evolved  by  some  ma- 
terials when  heated.  His  plan  was  to 
fill  an  elongated  glass  vessel  with 
mercury  (Fig.  19) ;  to  invert  this  in  a 
trough  filled  with  the  same  metal,  and, 
after  allowing  the  substance  under  ex- 
amination to  float  up  into  the  top  of  the 
tube  above  the  mercury,  to  expose  it  to 

the  rays  of  the  sun  concentrated  by  a  large  burning  lens.  Priestley 
found  that  one  material,  then  known  as  "  mercurius  calcinatus  per  se  " 
(mercuric  oxide),  gave  off  an  unusual  amount  of  a  gas,  or  "  air  "  as  he 
called  it.  He  found  that  this  gas  supported  combustion  extremely  well 
and,  later,  that  it  was  respirable  and  favorable  to  the  life  of  small  ani- 
mals, such  as  mice.  He  did  not,  however,  recognize  that  atmospheric 

61 


FIG.  19. 


62 


INORGANIC  CHEMISTRY 


air  was  a  mixture,  and  that  the  substance  he  had  obtained  was  in  reality 
identical  with  that  component  of  the  air  which  has  the  same  properties. 
Simultaneously  with  Priestley's  work,  Scheele  in  Sweden  had  been 
working  in  much  the  same  way,  and  had  obtained  the  same  gas  from 
niter,  mercuric  oxide,  and  other  substances,  publication  of  his  results 
being  delayed,  however,  until  1777. 

Meanwhile  Lavoisier,  in  Paris,  who  had  been  studying  the  rusting 
of  metals  in  the  air,  heard  of  Priestley's  experiments,  and  demonstrated 
that  the  latter's  "  good  air "  was  really  a  component  of  common  air, 
and  combined  with  metals  when  they  formed  rusts,  or  "  calces,"  as 
they  were  then  called.  He  proved  this  conclusively  by  heating  mer- 
cury in  the  retort  shown  in  Fig.  20.  The  apparatus  was  arranged  in 

such  a  way  that  a  definite  vol- 
ume of  air  was  inclosed,  within 
the  bell-jar  and  the  retort,  by 
the  larger  quantity  of  mercury 
which  filled  the  trough.  During 
the  heating  of  the  mercury  in 
the  retort,  the  familiar  red 
powder  (mercuric  oxide)  was 
formed  on  its  surface,  and  sim- 
ultaneously the  volume  of  air 
diminished,  until  after  twelve 
days  both  of  these  changes 
ceased.  The  air  had  suffered 

a  shrinkage  equal  to  about  one-fifth  of  its  volume,  and  an  easily  weigh- 
able  quantity  of  the  oxide  had  been  obtained.  The  gas  which  remained 
had  lost  the  power  of  supporting  combustion  and  life,  and  hence  was 
named  by  Lavoisier  "  azote  "  (Gk.  d  priv.  and  £urj,  life).  In  English  it 
is  called  nitrogen.  The  oxide  on  being  heated  more  strongly  by  itself 
gave  off  a  gas  whose  volume  exactly  corresponded  with  the  shrinkage 
undergone  by  the  inclosed  air,  and  this  gas  possessed  in  an  exaggerated 
degree  the  properties  which  the  air  had  lost.  The  proof  that  oxygen 
was  a  constituent  of  the  atmosphere  was  therefore  complete.  Lavoisier 
named  the  new  element  oxygen,  or  acid-producer  (Gk.  ££v?  an  acid,  ycwan 
to  produce),  from  the  fact  that  the  compounds  formed  by  its  union  with 
many  elements  gave  acid  (sour-tasting)  solutions  when  they  were  mixed 
with  water.  Cavendish  pointed  out  almost  immediately  that  there 
were  sour-tasting  substances  which  contained  no  oxygen,  so  that  the 
name  has  no  longer  any  significance. 


FIG.  20. 


OXYGEN  68 

Preparation  of  Simple  Substances.  —  There  are  two  general 
ways  of  obtaining  simple  substances.  If  the  element  occurs  uneom- 
bined  in  nature,  as  sulphur  and  gold  do,  it?  is  only  necessary  to  free  it 
from  foreign  materials  (impurities)  with  which  it  is  mixed.  If  no  such 
supply  exists,  or  if  the  purification  is  difficult,  then  some  compound, 
natural  or  artificial,  is  decomposed. 

The  decomposition,  in  turn,  may  be  effected  in  two  ways.  The 
compound  may  be  forced  apart  by  the  application  of  energy,  usually 
in  the  form  of  heat  or  electricity,  as  in  Priestley's  experiments  (cf.  pp. 
19-21).  Or  the  desired  constituent  may  be  liberated  by  offering  to  the 
other  constituents  some  substance  with  which  they  will  unite  (see 
Preparation  of  hydrogen).  When  oxygen  is  to  be  liberated,  the  former 
is  the  more  easily  applicable  plan. 

In  selecting  our  source,  we  are  naturally  influenced  by  the  cost  of 
the  material,  as  well  as  by  the  ease  of  the  process.  Thus,  gold  oxide 
yields  oxygen  by  the  application  of  very  little  heat,  but  it  is  extremely 
expensive.  Quicklime  is  very  cheap,  but  does  not  give  up  its  oxygen 
even  at  the  temperature  of  the  electric  arc. 

Preparation  of  Oxygen.  —  1.  Oxygen  may  be  separated  from 
the  other  substances  mixed  with  it  in  the  atmosphere  by  liquefying 
the  air  (see  Liquid  air),  allowing  the  nitrogen,  which  is  more  volatile, 
to  escape,  and  finally  compressing  into  tanks  the  oxygen  which  evap- 
orates last.  This  is  a  purely  mechanical  process. 

2.  There  are  many  compounds  which,  when  heated  to  temperatures 
under  2000°  such  as  we  can  obtain  with  the  aid  of  a  Bunsen  burner,  a 
coal  fire,  or  a  blast-lamp,  give  up  their  oxygen.  Some  of  them  are 
minerals,  but  most  of  them  are  manufactured  articles.  Of  the  min- 
erals, pyrolusite  (manganese  dioxide,  Mn02)  is  an  example.  It  usually 
contains  the  elements  of  water  also,  and  hence  moisture  is  evolved  at 
the  same  time.  A  substance  identical  with  the  mineral  hausmannite 
(Mn804)  remains.  Amongst  the  artificial  sources  are  mercuric  oxide, 
expensive,  but  historically  interesting  (p.  12) ;  barium  peroxide,  used 
in  manufacturing  oxygen  on  a  large  scale  (Erin's  process) ;  and  potas- 
sium chlorate,  the  most  convenient  for  laboratory  use.  Many  other 
substances  of  this  class  will  be  encountered  in  the  sequel. 

Brin's  Oxygen  Process  starts  from  barium  oxide  (g'.v.).  Barium 
oxide  BaO  closely  resembles  quicklime  CaO,  but  differs  from  this  sub- 
stance in  the  fact  that  when  heated  in  air  to  about  500°,  it  rapidly 
acquires  additional  oxygen  and  gives  barium  peroxide.  When  barium 


64  INORGANIC  CHEMISTRY 

peroxide  is  raised  to  a  higher  temperature  (1000°),  this  extra  oxygen 
is  given  up  again.  Barium  oxide  contains  one  chemical  unit  weight 
each  of  the  two  constituents  and  takes  up  another  of  oxygen,  so  that 
the  equation  for  the  primary  action  is : 

BaO  +  O  ->  Ba02. 

The  subsequent  decomposition  of  the  peroxide,  during  the  stage  in 
which  the  oxygen  is  made,  is  the  exact  opposite :  Ba02  — >  BaO  +  0.* 
The  commercial  advantage  of  the  method  lies  in  the  fact  that  the 
barium  oxide  remaining  after  the  second  stage  can  be  used  over  and 
over  again.  This,  as  will  be  seen,  is  in  reality  a  chemical  method  of 
obtaining  oxygen  from  the  air. 

In  practice  an  improvement  on  the  above  principle  makes  working 
more  economical.  It  is  found  that  if  the  barium  oxide  is  maintained 
at  a  temperature  of  700°,  intermediate  between  the  two  just  mentioned, 
oxygen  is  absorbed  when  air  is  forced  under  pressure  into  the  tubes 
containing  the  oxide.  A  valve  at  the  extremity  of  the  tubes  permits 
the  escape  of  the  nitrogen.  When  the  combination  with  oxygen  is 
completed,  the  pumping  apparatus  is  reversed,  and,  a  partial  vacuum 
being  created,  the  oxygen  in  combination  is  given  off  without  any 
alteration  in  temperature  being  necessary.  Thus  a  great  waste  of  fuel 
is  avoided,  and  the  process  is  rendered  more  nearly  continuous.  This 
method  furnishes  oxygen  about  96  per  cent  pure,  and  suitable  for  sale 
in  compressed  form  in  cylinders. 

Potassium  Chlorate  (q.v.)  is  a  white  crystalline  substance  used  in 
large  quantities  in  the  manufacture  of  matches  and  fireworks.  When 
heated  to  a  moderately  high  temperature  in  a  tube  similar  to  that  in 
Fig.  5,  it  gives  off  a  very  large  volume  of  oxygen.  Examination  shows 
that  the  whole  of  the  oxygen  it  contains  can  be  driven  out.  The  white 
material  which  remains  after  the  heating  is  identical  with  the  mineral 
sylvite.  To  the  chemist  it  is  known  as  potassium  chloride,  and,  when 
decomposed,  it  yields  one  atomic  weight  each  of  potassium  and  chlo- 
rine. Its  formula  is  thus  KC1.  We  may  infer,  therefore,  that  the 
composition  of  the  original  substance  will  be  representable  by  the  for- 
mula KC1OZ,  where  x  is  the  number  of  atomic  weights  of  oxygen. 
Measurement  and  calculation  show  x  =  3.  The  formula  is  therefore 

*  In  cases  where  an  action  is  reversible,  and  the  direction  depends  on  condi- 
tions which  may  be  altered,  we  write  both  equations  in  one : 

BaO  +  O  ^±  Ba02. 


OXYGEN 


65 


KC103,  and  the  equation  for  the  decomposition  (see,  however,  under 
Perchlorates,  p.  275)  : 

KC108  ->  KC1  +  3O. 

To  learn  the  value  of  x,  we  ascertain  the  loss  in  weight  ( =  oxygen)  which  a 
known  quantity  of  potassium  chlorate  sustains  when  heated  in  a  hard  glass  tube 
closed  at  one  end.  By  subtraction  we  get  the  weight  of  potassium  chloride  for- 
merly combined  with  the  oxygen.  In  an  actual  experiment,  2.998  g.  of  potassium 
chlorate  gave  1.169g.  of  oxygen  and  left  1.829  g.  of  the  chloride.  The  atomic 
weights  of  potassium  and  chlorine  are  39.15  and  35.45  respectively,  and  the 
formula-weight  of  the  chloride  is  therefore  74.6.  Dividing  the  measured  weights 
of  oxygen  and  potassium  chloride  by  the  corresponding  atomic  and  formula  weights 
(cf.  p.  67)  respectively:  1.169  -^  16  =  0.07306  and  1.829  ~  74.6  =  0.02452.  We 
observe  that  the  ratio  of  the  quotients  is  2.98  :  1,  or  almost  exactly  3 :  1.  The 
formula  is  therefore  O  x  3,  (KC1)  x  1,  or  KC1O8. 

A  peculiarity  of  this  action  is  that  admixture  of  manganese  dioxide 
increases  very  markedly  the  speed  with  which  the  decomposition  of 
the  potassium  chlorate  takes  place.  Hence,  in  its  presence,  and  it  is 


FIG.  21. 

generally  mixed  with  the  chlorate  in  laboratory  experiments  (Fig.  21), 
a  sufficient  stream  of  the  gas  is  obtained  at  a  relatively  low  tempera- 
ture (below  200°).  Without  it  no  oxygen  is  evolved  at  all  until  the 
chlorate  melts  (351°).  The  dioxide  does  not  begin  to  lose  oxygen 
below  400°,  and  therefore  is  not  itself  permanently  changed  in  any 
way  when  used  for  this  purpose. 

Familiarity  with  Physics  Required  in  the  Study  of  Chem- 
istry.—  In  mentioning  chemical  phenomena,  it  is  inevitable  that  considerations 
of  space  should  limit  our  statements  to  the  merest  indication  of  the  process  and 
the  briefest  record  of  the  chemical  result.  The  prodigious  disproportion  between 


66  INORGANIC  CHEMISTRY 

the  meagerness  of  this  fragment  and  the  mass  of  detail  which  lies  behind  it  in  each 
case  should  be  constantly  before  the  mind  of  the  reader.  The  book  gives  empiri- 
cal knowledge,  the  laboratory  work  and  the  discussion  of  it  furnish  the  only  real 
knowledge.  The  extent  and  nature  of  this  real  knowledge  may  be  shown  in  con- 
nection with  any  action.  As  an  illustration  we  may  point  out  some  of  the  prob- 
lems which  the  heating  of  potassium  chlorate  presents  to  one  who  is  trying  to 
acquire  an  intelligent  acquaintanceship  with  its  chemistry,  and  has  not  previously 
done  the  experiment. 

First,  the  substance  melts.  It  must  be  realized  that  this  is  a  common  occur- 
rence which  does  not  necessarily  imply  any  profound  change,  and  may  be  reversed 
by  cooling.  Later,  the  liquid  appears  to  boil,  and  the  properties  of  a  boiling  sub- 
stance must  be  known.  If  the  observer  has  been  informed  in  advance  that  the 
body  is  homogeneous,  he  must  know  that,  if  it  is  simply  boiling,  it  will  evaporate 
completely  and  leave  nothing  behind,  and  that  the  temperature  required  to  achieve 
this  will  remain  constant  from  the  beginning  to  the  end.  In  order,  therefore,  to 
become  aware  of  the  fact  that  here  decomposition  is  taking  place,  he  must  note  the 
ways  in  which  the  decomposition  of  potassium  chlorate  differs  from  ordinary  boil- 
ing. For  example,  if  it  were  a  case  of  boiling,  he  should  expect  to  find  the  solid 
body  condensing  on  the  sides  of  the  tubes,  and  note  the  fact  that  no  such  conden- 
sation is  observed,  with  the  appropriate  inferences.  He  should  observe  that,  in 
the  later  stages  at  least,  the  agitation  of  the  liquid  does  not  cease  when  the  flame 
is  removed,  although  this  would  undoubtedly  occur  in  a  case  of  simple  boiling. 
He  must  further  observe  the  changes  in  the  consistency  of  the  material  and  the 
way  in  which  it  finally  becomes  thick  and  may  even  solidify.  Even  the  most  ex- 
perienced investigator  would  have  to  make  many  careful  experiments  before  he 
could  definitely  classify  the  nature  of  the  phenomenon  being  observed.  The  first 
inference  would  probably  be  that  the  phenomenon  was  certainly  not  one  of  mere 
ebullition.  In  some  ways  it  is  like  the  evaporation  of  a  solution,  obtained,  say, 
by  the  melting  of  a  substance  in  its  own  water  of  crystallization.  Yet,  this  hypoth- 
esis would  not  explain  to  the  thoughtful  observer  even  the  more  obvious  features 
of  the  phenomenon,  for  the  liquid  which  was  acting  as  a  solvent  would  have  to  be 
amazingly  volatile  if  the  absence  of  any  condensation  on  the  walls  of  the  tube  was 
to  be  accounted  for. 

The  illustration  need  not  be  elaborated  further.  These  remarks  are  sufficient  to 
show  that  even  the  simplest  experiment  presents  an  almost  limitless  field  for  the 
discussion  of  important  questions  which  are  more  or  less  common  to  all  chemical 
phenomena.  It  must  be  noted,  also,  that  chemical  change  is  in  itself  not  percept- 
ible by  the  senses,  and  that  only  physical  properties  and  physical  phenomena  are 
observed  (cf.  pp.  36-40).  The  chemical  facts,  such  as  the  general  nature  of  the 
change,  the  conditions  under  which,  and  the  facility  with  which  it  occurs,  are 
reached  solely  by  inference.  The  above  example  shows  the  ready  and  thorough 
knowledge  of  physics  which  must  be  at  the  command  of  every  individual  effort  to 
study  even  the  simplest  chemical  phenomenon.  It  is  only  when  the  physics  as 
well  as  the  chemistry  of  the  change  have  been  mastered  that  the  "  real  knowledge  " 
to  which  reference  was  made  above  has  been  gained. 

Physical  Properties  of  Oxygen.  —  Oxygen,  as  a  gas,  resembles 
air  in  being  colorless,  tasteless,  and  odorless.  It  is  slightly  heavier 


OXYGEN  67 

than  air ;  its  density,  using  the  physical  standard  of  air  =  1,  is  1.105. 
The  chemist  often  uses  hydrogen  as  his  standard,  and  oxygen  is 
15.900  times  (Morley)  as  heavy.  One  liter  of  oxygen,  at  0°  and  760 
mm.  barometric  pressure,  weighs  1.42900  grams  (Morley).  The  gas 
dissolves  to  some  extent  in  water,  the  solubility  at  0D  being  four  vol- 
umes of  gas  in  one  hundred  volumes  of  water  (at  20°,  3 : 100).  The 
critical  temperature  (q.v.)  of  oxygen  is  — 118°  C.  At  that  tempera- 
ture, fifty  atmospheres  pressure  is  required  to  liquefy  it.  Liquid 
oxygen,  which  was  first  made  by  Wroblewski,  has  a  pale-blue  color, 
and  boils  under  one  atmosphere  at  — 182.5°.  Its  specific  gravity  at 
-182.5°  is  1.13  (water  =  1) :  that  is  to  say,  1  c.c.  weighs  1.13  g.  By 
cooling  with  a  jet  of  liquid  hydrogen,  Dewar  froze  the  liquid  to  a 
snow-like,  pale-bluish  solid.  The  magnetic  properties  of  the  element 
are  clearly  shown  by  the  force  with  which  a  tube  of  liquid  oxygen  is 
attracted  by  a  magnet. 

Specific  Chemical  Properties.  —  Under  this  head  we  describe 
the  chemical  behavior  of  a  substance,  enumerating  the  other  sub- 
stances, simple  or  compound,  with  which  it  unites  or  interacts  (p.  14) , 
stating  the  conditions  peculiar  to  each  action,  and  estimating  the 
intensity  of  the  tendency  to  chemical  change  in  each  case.  In  the 
case  of  a  simple  substance  like  oxygen  we  note  particularly  with  how 
many  of  the  other  elements  it  can  form  compounds,  how  far  it  unites 
with  them  directly,  and  in  how  many  cases  the  compounds  have  to  be 
made  by  indirect  means.  In  general,  we  call  those  simple  substances 
active  which  unite  with  many  other  simple  substances  and  do  so  by 
direct  union.  Oxygen,  for  example,  is  active,  and  nitrogen  (q.v.)  is 
relatively  inert.  The  intensity  of  a  chemical  action  is  judged  by  its 
speed,  by  the  electricity  it  can  generate,  or,  roughly,  by  the  heat  evolved 
during  its  progress  (p.  28). 

Chemical  Properties  of  Owygen.  —  Sulphur,  when  raised  in 
advance  to  the  temperature  necessary  to  start  the  action,  unites  vigor- 
ously with  oxygen  (Fig.  22),  giving  out  much  heat  and  producing  a 
familiar  gas  having  a  pungent  odor  (sulphur  dioxide).  This  odor  is 
frequently  spoken  of  as  the  "  smell  of  sulphur,"  but  in  reality  sulphur 
itself  has  no  odor,  and  neither  has  oxygen.  The  odor  is  peculiar  to 
the  compound  of  the  two.  It  may  be  added  that  the  mode  of  experi- 
mentation can  be  changed  and  the  oxygen  led  into  sulphur  vapor 
through  a  tube.  The  former  then  appears  to  burn  with  a  bright 
flame,  giving  the  same  product  as  before. 


68 


INORGANIC   CHEMISTRY 


Warm  phosphorus  combines  with  oxygen  with  even  greater  vigor, 
and  forms  a  white,  powdery,  solid  compound  (phosphoric  anhydride), 
which  absorbs  moisture  from  the  aqueous  vapor  in  the  air  and  quickly 
forms  a  solution  in  this  water.  In  both  these  cases  the  products  differ 

from  oxygen,  not  only  in  odor  and 
in  being  a  gas  and  a  solid,  respec- 
tively, but  notably  in  that,  when 
shaken  with  water,  they  dissolve 
and  interact  to  form  acids  (see 
below). 

Burning  carbon,  in  the  form  of 
charcoal,  when  plunged  into  the  gas, 
glows  much  more  brightly  than  in 
ordinary  air.  In  this  case  the  pro- 
duct is  a  gas  (carbon  dioxide). 
When  this  gas  is  shaken  with 
"  lime-water,"  a  solution  of  calcium 
hydroxide  Ca(OH)2  (q.v.),  a  white 
precipitate  of  calcium  carbonate 
CaCOg  is  formed. 

Finally,  metallic  iron,  which  is 
simply  rusted  by  air  (diluted  oxy- 
gen), burns  in  pure  oxygen  with 
surprising  brilliancy.  Globules  of 
a  molten  product  fall  from  the  iron, 
and  when  they  have  cooled  are 
found  to  consist  of  a  dark-gray 

brittle  material,  which  we  recognize  as  identical  with  blacksmith's 
hammer-scale,  and  with  a  well-known  ore  of  iron  (magnetic  oxide  of 
iron). 

Experiments  like  the  above  with  the  whole  series  of  simple  sub- 
stances would  show  that  oxygen  unites  in  a  similar  way  with  nearly 
every  member  of  the  list,  and  often,  though  not  always,  with  the  same 
vigor  as  in  the  case  of  these  examples.  In  the  case  of  one  or  two 
elements,  such  as  gold  and  platinum,  the  compounds  are  obtainable  by 
double  decomposition,  and  not  directly.  With  the  five  members  of  the 
helium  group,  of  which  no  chemical  compounds  are  known,  and  with 
fluorine,  oxygen  does  not  combine.  These  three  sentences  summarize 
the  chemical  properties  of  oxygen. 

Oxygen  can  unite  with  many  of  the  same  elements  when  they  are 


FIG.  22. 


OXYGEN  69 

already  in  combination.  Wood,  for  example,  is  composed  of  carbon 
and  hydrogen,  with  a  certain  amount  of  oxygen.  When  previously 
heated,  it  is  decomposed,  and  the  constituents  unite  with  oxygen  form- 
ing carbon  dioxide  and  water. 

The  Making  of  Equations  Again.  —  To  learn  the  exact  nature 

of  interactions  like  those  used  as  illustrations  above,  quantitative  ex- 
periments must  of  course  be  made  and  the  usual  method  employed  to 
obtain  the  formulae  of   the  products.     Thus,  for  example,  a  known 
weight    of      sul- 
phur is  placed  in 

a  porcelain  boat         c=il_E ^  2 • 

(Fig.  23),  which 
has  already  been 
weighed.  The 
U-shaped  tube  to 
the  right  con- 
tains a  solution  FlGt  23- 
of  potassium  hy- 
droxide which  is  capable  of  absorbing  the  resulting  gas.  The  oxygen 
enters  from  the  left.  When  the  sulphur  is  heated,  it  burns  in  the 
oxygen,  and  the  loss  in  weight  which  the  boat  undergoes  shows  the 
amount  of  sulphur  consumed.  The  gain  in  weight  of  the  U-tube 
shows  the  weight  of  the  compound  produced.  By  subtracting  from 
this  weight  that  of  the  sulphur,  we  get  the  quantity  of  oxygen.  The 
proportion  of  the  constituents  and  the  steps  in  the  calculation  are  as 
follows : 

PERCENTAGE*          AT.  WT.        FACTOR  -^- 1.561 

Sulphur,          50.05        =        32.06     x     1.661       =      S  x  1.561         S  x  1 
Oxygen,          49.95        =        16.00     x     3.122      =      Ox  3.122        O  x  2 

The  formula  of  the  product  is  therefore  SO2  and  the  equation  S  -f  20 
->S02. 

Similarly,  phosphoric  anhydride  may  be  shown  to  have  the  for- 
mula P205,  carbon  dioxide  C02,  and  magnetic  oxide  of  iron  Fe804. 

The  results  given  by  the  above  experiment  are  usually  inexact.  The  tendency 
to  the  formation  of  sulphur  trioxide,  often  heightened  by  catalytic  action  (see 

*  Here  the  percentages  are  employed.  The  actual  weights  found  in  one  experi- 
ment, however,  may  be  divided  by  the  atomic  weights  and  the  same  result  obtained 
(p.  65) .  In  fact,  any  two  numbers  bearing  the  proper  ratio  to  one  another  may  be 
used- 


70  INORGANIC   CHEMISTRY 

below)  of  the  porcelain  of  the  boat,  raises  abnormally  the  proportion  of  oxygen. 
The  principle  of  the  experiment  is  easy  to  understand,  however. 

In  the  case  of  phosphorus  a  similar  plan  of  experiment  may  be  used.  Instead  of 
attempting  to  receive  the  solid  product  in  a  U-tube,  however,  it  must  be  caught  by 
a  plug  of  glass  wool  in  the  main  tube  of  hard  glass,  and  a  drying  tube  will  be 
needed  at  the  end  to  prevent  admission  of  moisture  from  the  air.  The  increase  in 
weight  of  the  hard  glass  tube  represents  the  oxygen  taken  up.  Care  and  leisurely 
performance  are  needed  to  make  the  experiment  successful. 

The  measurement  of  the  composition  of  carbon  dioxide  *  gives  the  most  exact 
results  when  carefully  conducted.  The  data  and  working  in  these  cases  and  in 
that  of  iron  are  as  follows  : 

PERCENTAGE  AT.  WT.  FACTOR  -f-  .704 

Phosphorus,  43.66  31.0  1.408  P  x  2 

Oxygen,  56.34  16.0          x          3.521  O  x  5 

-=-2.272 

Carbon,  27.27          =          12.0          x          2.272  C  x  1 

Oxygen,  72.72  16.0         x          4.545  O  x  2 

-=-0.431 

Iron,  72.38          =          55.9          x          1.295  Fe  x  3 

Oxygen,  27.62          -          16.0          x          1.726  O    x  4 

The  density  of  the  vapor  of  phosphoric  anhydride  leads  to  the  formula  P4O10, 
but,  as  the  substance  is  never  used  as  a  gas,  the  simpler  formula  is  generally 
preferred. 

Oxides.  —  Substances  containing  one  element  in  combination  with 
oxygen  are  called  oxides,  and  processes  like  those  described  above  are 
called  oxidizing  processes,  or  oxidations.  When  the  same  element 
forms  more  than  one  oxide,  the  names  of  the  oxides  indicate  the  differ- 
ing proportions.  Thus  we  have  barium  oxide  (or  monoxide)  BaO,  and 
barium  peroxide  (or  dioxide)  Ba02,  magnetic  oxide  of  iron  Fe304,  fer- 
rous oxide  FeO,  and  ferric  oxide  Fe203.  In  cases  like  the  two  last 
the  terminations  -ous  and  -ie  applied  to  the  metal  correspond  to  the 
smaller  and  larger  proportions  of  oxygen,  respectively,  which  the  metal 
is  able  to  hold  in  combination.  Oxides  of  the  form  Fe203  are  often 
called  sesquioxides  because  they  contain  one  and  a  half  (Lat.  sesqui-, 
one-half  more)  unit  weights  of  oxygen  to  each  one  of  iron. 

Many  oxides,  like  those  of  iron,  are  quite  indifferent  to  water,  but 
others,  like  those  of  sulphur  and  phosphorus,  interact  with  it  (see 
under  Water).  Some  give  sour  solutions,  containing  acids  dissolved  in 
the  excess  of  water.  Such  solutions  turn  blue  litmus,  a  vegetable  dye, 
red.  Others  give  solutions  with  a  taste  like  soap  or  borax,  and  here 

*  Described  in  the  author's  Laboratory  Outline  of  General  Chemistry. 


OXYGEN  71 

the  dissolved  substance  is  called  a  base  (q.v.),  and  turns  litmus  blue. 
Thus  sulphur  dioxide  and  phosphoric  anhydride  give  sulphurous  acid 
and  phosphoric  acid  respectively  : 

S02  +  H20  ->  H2S08 
P206  +  3H20  -,  2H8P04.* 

If  the  product,  to  whichever  class  it  belongs,  is  not  volatile,  it  may  be 
obtained  by  evaporating  the  excess  of  water.  In  the  case  of  sulphur- 
ous acid  the  above  action  is  reversed  by  evaporation  and  the  sulphur 
dioxide  and  water  both  pass  off ;  in  that  of  phosphoric  acid,  the  white 
crystalline  acid  is  obtained.  In  consequence  of  their  relation  to  the 
acid,  differing  from  it  in  not  containing  the  elements  of  water,  these 
oxides  are  often  called  anhydrides. 

The  discussion  of  the  formation  and  properties  of  ozone  (q.v.), 
which  is  an  oxide  of  oxygen,  cannot  be  taken  up  until  we  are  in  pos- 
session of  the  means  of  understanding  the  difference  between  the  two 
substances. 

Combustion.  —  Since  oxygen  is  a  component  of  the  atmosphere, 
chemical  actions  in  which  it  plays  a  part  are  familiar  in  daily  life. 
Violent  union  with  oxygen  is  called  in  popular  language  combustion 
or  burning.  Yet  since  oxygen  is  only  one  of  many  gaseous  substances 
known  to  the  chemist,  and  similar  vigorous  interactions  with  these 
gases  are  common,  the  term  has  no  scientific  significance.  The  union 
of  iron  and  sulphur,  even,  gives  out  light  and  heat,  and  is  quite  simi- 
lar in  the  chemical  point  of  view  to  combustion. 

In  connection  with  this,  however,  it  may  be  worth  while  to  notice 
the  distinction  between  combustible  and  incombustible  substances. 
Things  which  are  incombustible,  using  the  term  in  its  popular  sense, 
may  be  divided  into  two  classes.  There  are  those  substances  which 
already  contain  all  the  oxygen  which  they  can  hold  in  combination. 
Such  are  the  oxides  whose  formation  we  observed  in  the  experiments 
described  above.  In  everyday  life,  limestone,  sand,  bricks,  and  most 
rocks  are  illustrations.  The  other  substances  ordinarily  classed  as 
incombustible  are  those  which  do  not  unite  with  oxygen,  as  it  is  found 
in  the  air,  with  sufficient  vigor.  The  iron  used  in  the  construction  of 
fireproof  buildings  is  the  commonest  example  of  this  class. 

*  Here  summation  of  the  formulae  on  the  left  would  give  H6P2O8,  but  in  such 
cases,  unless  there  are  reasons  to  the  contrary,  the  common  factor  is  put  in  front 
and  the  formula  reduced  to  its  lowest  terms. 


72  INORGANIC   CHEMISTRY 

Oxidation.  —  The  rusting  of  metals  differs  from  combustion  only 
in  speed.  Thus,  magnesium  ribbon,  when  left  lying  in  the  air,  gradu- 
ally becomes  covered  with  a  white  crust.  By  scraping  this  off,  and 
continually  exposing  the  surface  of  the  metal,  the  whole  of  the  latter 
may  eventually  be  transformed  into  the  white  powder.  The  product 
(provided  carbon  dioxide  was  not  present)  is  the  oxide  formed  by 
combustion  (p.  44).  In  the  case  of  iron,  burning  gives  us  the 
magnetic -oxide  (Fe304),  while  rusting  in  moist  air  yields  a  hydrated 
ferric  oxide  (Fe2O3  4-  Aq*).  The  products  differ  in  composition,  but 
are  closely  related. 

This  process  of  slow  oxidation,  although  less  conspicuous  than 
combustion,  is  really  of  greater  interest.  Thus  the  decay  of  wood  is 
simply  a  process  of  oxidation  whereby  the  same  products  are  formed 
as  by  the  more  rapid  ordinary  combustion.  Large  volumes  of  pure 
water  are  mixed  with  sewage,  the  object  being  not  simply  to  dilute  the 
latter  but  to  introduce  water  containing  oxygen  in  solution.  This  has 
an  oxidizing  power  like  that  of  oxygen  gas,  and,  through  the  agency  of 
bacteria,  quickly  renders  dissolved  organic  matters  innocuous  by  con- 
verting them  for  the  most  part  into  carbon  dioxide  and  water.  In  our 
own  bodies  we  have  likewise  a  familiar  illustration  of  slow  oxidation. 
Avoiding  details,  it  is  sufficient  to  say  that  the  oxygen  from  the  air 
taken  into  the  lungs  is  carried  by  the  blood  throughout  our  tissues  and 
is  there  used  for  oxidizing  waste  materials.  The  carbon  dioxide  is 
carried  back  to  the  lungs  by  the  blood,  and  finally  reaches  the  air  dur- 
ing exhalation.  To  supply  the  place  of  the  material  thus  removed,  we 
are  under  the  continual  necessity  of  building  new  tissue  from  the  food 
which  we  eat.  If  we  cease  to  eat,  we  become  lighter  and  weaker,  show- 
ing that  a  real  portion  of  our  structure  is  gradually  being  consumed  by 
oxidation. 

The  opposite  of  oxidation,  the  removal  of  oxygen,  is  spoken  of  in 
chemistry  as  reduction. 

Means  of  Altering  the  Speed  of  a  Criven  Chemical  Action :  By 
Change  of  Temperature.  —  That  the  same  change  may  proceed  with 
very  different  speeds  according  to  conditions  is  a  familiar  fact.  For 
example,  raising  the  temperature  increases  the  rapidity  of  all  chemical 
interactions.  Thus,  cold  iron  combines  with  oxygen  very  slowly,  giv- 

*  The  formula  H20  may  not  be  used  excepting  to  indicate  a  definite  proportion 
of  the  elements  of  water  (18  parts).  Where  the  proportion  varies  according  to 
circumstances,  as  here  and  in  the  case  of  solutions,  the  contraction  Aq  is  employed. 


OXYGEN  73 

ing  rust,  while  white-hot  iron  sheds  quantities  of  scales  of  an  oxide, 
formed  in  the  few  moments  that  it  is  under  the  blacksmith's  hammer. 
White-hot  coal  unites  with  oxygen  in  the  air  to  form  carbon  dioxide 
and  seems  to  disappear  before  our  eyes,  while  in  the  cellar,  even  in 
warm  weather,  we  observe  no  appreciable  diminution  in  its  amount. 
The  chemist,  however,  has  reason  for  considering  that  even  here  the 
difference  is  one  of  degree  only.  No  temperature  can  be  found  at  which 
the  interaction  definitely  begins.  We  believe  that  every  change,  pro- 
vided, like  these,  it  involves  a  liberation  of  energy  (cf.  p.  27),  proceeds 
with  some  speed  at  every  temperature.  A  rough  estimation,  based  on 
experiment,  shows  that  on  an  average,  other  things  being  equal,  every 
rise  in  temperature  of  ten  degrees  doubles  the  amount  of  material 
changed  per  second,  and  conversely. 

If,  on  bringing  two  materials  together,  the  chemist  observes  no 
marks  of  chemical  action,  he  immediately  begins  cautiously  to  heat  the 
mixture.  This  appeal  to  the  accelerating  effect  of  a  rise  in  tempera- 
ture is  always  made  as  a  matter  of  course. 

The  common  expressions  used  in  chemistry  in  describing  temper- 
atures, along  with  the  corresponding  readings  of  the  thermometer,  are 
as  follows  : 

Incipient  red  heat,  about  525°.     Yellow  heat,  about  1100°. 

Dark  red  heat,  "    700°.     Beginning  white  heat,  «  1300°. 

Bright  red  heat,  "    950°.     White  heat,  "  1500°. 

Rapid  Self-sustaining  Chemical  Action  and  Means  of  Initiat- 
ing it.  —  When  a  piece  of  wood  is  set  on  fire  at  one  end,  the  heat  pro- 
duced by  the  action  itself  raises  the  temperature  of  neighboring  portions 
until  their  speed  of  union  becomes  equal  to  that  of  the  part  originally 
lighted.  In .  this  way  the  whole  becomes  finally  inflamed.  When  we 
blow  the  blaze  out,  the  great  excess  of  cold  air  suddenly  lowers  the 
temperature  of  the  wood,  and  of  the  gas  rising  from  it,  and  rapid  union 
ceases.  Whether  a  given  set  of  materials  can  maintain  itself  at  a  tem- 
perature proper  to  violent  interaction  will  depend  on  the  amount  of  heat 
developed  by  the  action  itself  on  the  one  hand,  and  the  losses  of  heat 
by  conduction  and  radiation  on  the  other.  If  the  latter  are  great,  the 
former  must  be  greater.  Thus  the  union  of  iron  and  oxygen  per  se 
gives  heat  enough  to  warm  the  materials  to  the  burning  temperature  and 
leaves  much  over  for  radiation.  But  iron  in  air,  which  is  four-fifths 
nitrogen,  can  receive  the  oxygen  only  one-fifth  as  fast  at  the  start,  and 
even  more  slowly  as,  later,  the  nitrogen  accumulates  round  it.  And  be- 


74  INORGANIC  CHEMISTRY 

sides,  all  the  nitrogen  has  to  be  heated  to,  perhaps,  2000°.  The  task  is 
too  great.  The  union  is  impeded  and  the  iron  is  not  oxidized  fast 
enough  to  generate  the  heat  required  to  maintain  everything  at  this 
high  temperature.  Poor  conductors  of  heat,  like  wood  and  candles,  fare 
better.  Powdered  iron,  with  its  particles  presenting  large  surface  to 
the  air  relatively  to  the  weight  of  material  in  each  particle  to  be  heated, 
burns  well. 

The  initial  supply  of  heat  required  to  start  violent  exothermal  chem- 
ical actions,  of  which  alone  we  are  here  speaking,  must  not  be  confused 
with  the  heat  subsequently  developed  as  the  action  proceeds.  The  lat- 
ter is  usually  much  greater.  Indeed,  the  preliminary  supply  varies 
with  circumstances,  and  may  be  made  as  small  as  we  choose  by  limiting 
the  area  first  heated  and  using  ordinary  precautions  against  radiation 
and  convection.  Indeed,  in  practice,  a  single  spark  from  an  induction 
coil  often  takes  the  place  of  more  clumsy  methods  of  raising  the  tem- 
perature. The  heat  produced  by  the  interaction  itself,  however,  is  fixed 
in  amount,  and  depends  only  on  the  materials  and  their  quantity.  En- 
dothermal  actions  differ  completely  from  those  under  discussion.  In 
them  a  definite,  large  quantity  of  heat  has  to  be  furnished,  and  the  ac- 
tion instantly  ceases  if  the  supply  fails. 

Heating  is  not  the  only  means  used  to  give  the  initial  acceleration 
to  a  self-sustaining  chemical  change.  The  materials  in  a  match-head 
are  capable  of  undergoing  a  great  transformation.  Yet,  so  slowly  does 
this  proceed  at  ordinary  temperatures,  that  matches  may  be  kept  in 
efficient  condition  for  years.  Here  a  rather  violent  vibration  is  em- 
ployed to  hasten  the  torpid  action  in  a  small  part  of  the  material,  and 
the  heat  produced  by  the  resulting  action  quickly  ignites  the  whole. 
The  same  explanation  accounts  for  the  explosion  of  gun-cotton  by  a 
percussion  fuse. 


Other  Means  of  Altering  the  Speed  of  a  Given  Chemical 
Change :  By  Change  in  Concentration  ;  by  Catalysis  ;  by  Solu- 
tion. —  Even  when  the  temperature  remains  constant,  there  are  other 
changes  in  the  conditions  (p.  52)  which  may  be  used  for  accelerating  or 
for  moderating  the  speed  of  chemical  interactions.  The  most  impor- 
tant of  these  is,  a  change  in  the  concentration  of  the  interacting  sub- 
stances. Another  is  the  presence  of  a  catalytic  agent.  The  condition 
of  solution  might  be  accounted  still  another. 

The  abatement  in  the  activity  of  the  oxygen  found  in  the  air  by  the 
nitrogen  which  is  mixed  with  it,  is  a  question  of  concentration  If  the 


OXYGEN  75 

concentration  of  pure  oxygen  under  atmospheric  pressure  is  taken  as 
unity,  that  of  oxygen  in  air  is  only  about  0.2.  And  the  speed  of  inter- 
action of  a  body,  other  things  being  equal,  is  directly  proportional  to 
its  concentration.  This  is  not  an  obscure  law,  but  merely  common 
sense  put  into  definite  language.  The  opportunity  which  one  substance 
has  for  getting  at  every  part  of  another  will  be  one  factor  in  determin- 
ing the  speed  with  which  the  resulting  transformation  will  take  place. 
And  this  opportunity,  other  things  being  equal,  depends  on  the  thick- 
ness or  density  with  which  the  substance  is  scattered  in  the  region  of 
action.  In  the  case  of  a  gas,  this  factor  is  measured  by  its  partial  press- 
ure. Hence,  lights  burn  badly  at  great  elevations,  where  the  oxygen 
is  very  tenuous.  On  the  other  hand,  powdered  charcoal  interacts  so 
rapidly  when  ignited  in  liquid  air,  where  the  oxygen  is  highly  con- 
densed, that  an  explosion  takes  place,  whereas  in  common  air  it  burns 
feebly.  Again,  when  oxygen  is  compressed  in  contact  with  barium 
oxide  at  700°  it  combines  to  form  the  dioxide  (p.  64)  ;  when  the  press- 
ure of  the  oxygen  in  contact  with  the  latter  is  reduced,  oxygen  is  lib- 
erated (see  Chemical  equilibrium). 

When  an  extra  substance  increases  the  speed  of  a  chemical  change, 
seemingly  by  its  mere  presence,  without  itself  suffering  any  permanent 
change,  we  call  this  catalytic  (Gk.  Kara.,  down ;  Awns,  the  act  of  loosing) 
or  contact  action.  The  word  was  originally  used  for  cases  of  decomposi- 
tion. The  foreign  body  is  called  the  catalytic  or  contact  agent,  and  the 
process  catalysis.  The  effect  of  manganese  dioxide  on  the  decompo- 
sition of  potassium  chlorate  (p.  65)  is  of  this  nature.  When  some  of 
the  chlorate  is  melted  carefully,  so  as  to  avoid  superheating,  scarcely 
any  evolution  of  oxygen  can  be  perceived  at  this  temperature  (351°). 
If  now  a  thin  glass  bulb  filled  with  powdered  manganese  dioxide  be 
broken  in  the  molten  mass,  the  oxygen  is  given  off  in  torrents  in  con- 
sequence of  the  enormous  acceleration  of  the  decomposition.  Yet  the 
manganese  dioxide  itself  may  be  recovered  unchanged  from  the  residue. 

It  is  found  that  many  actions  owe  what  appears  to  be  their  normal 
speed  to  the  presence  of  a  trace  of  water  vapor.  Thus  many  of  the 
elements  show  no  visible  tendency  to  unite  with  carefully  dried  oxy- 
gen, even  when  they  are  strongly  heated  in  it.  Addition  of  a  trace  of 
moisture,  however,  brings  about  instant  combustion.  So  water  is  to  be 
regarded  as  one  of  the  commonest  contact  agents  (see  Chemical  prop- 
erties of  chlorine). 

A  few  cases  of  retardation  of  an  action  by  a  catalytic  agent  are  known.  Thus, 
a  little  benzyl  alcohol  or  mannite  added  to  the  solution  will  retard  the  oxidation  of 


76  INORGANIC  CHEMISTRY 

sulphites  by  the  air.  Hence  positive  and  negative  catalysis  both  occur  (see  ab- 
stract of  an  address  by  Ostwald  in  Nature,  LXV,  622,  given  also  in  full  in  the 
Zeitschriftfiir  Elektrochemie  VII,  995). 

The  effect  of  solution  in  hastening  a  chemical  change  was  seen 
when  we  examined  the  interaction  of  sodium  chloride  and  silver  nitrate 
(p.  13).  With  the  solutions  the  action  was  seemingly  instantaneous. 
If  we  had  attempted  to  bring  it  about  by  rubbing  the  dry  substances  in 
a  mortar,  hours  of  work  would  have  left  much  of  the  original  bodies 
still  unchanged.  Even  heating  would  not  have  produced  so  prompt  an 
effect.  It  is  obvious  that  the  intimate  access  which  every  part  of  each 
solution  gains  to  every  part  of  the  other  accounts  to  some  extent  for 
the  difference  (see  lonization).  Chemical  actions,  as  will  be  seen  in  the 
sequel,  are  very  frequently  carried  out  in  aqueous  solution  in  order  to 
take  advantage  of  the  favorable  influence  of  this  condition. 

Thermochemistry.  —  As  we  have  seen  (p.  27),  the  free  or  available 
chemical  energy  of  a  system  undergoing  chemical  change  usually  ap- 
pears in  the  form  of  heat.  Since  it  is  often  instructive  to  consider  the 
amount  of  heat  produced,  some  of  the  elementary  facts  of  thermo- 
chemistry must  be  explained. 

The  chemical  interactions  to  be  studied  thermally  are  arranged  so 
that  they  may  be  carried  out  in  some  small  vessel  which  can  be  placed 
inside  another  containing  water.  The  heat  developed  raises  the  tem- 
perature of  this  water.  Where  gases  like  oxygen  are  concerned,  a 
closed  bulb  of  platinum  forms  the  inner  vessel.  The  quantity  of  heat 
capable  of  raising  one  gram  of  water  one  degree  in  temperature,  be- 
tween 0°  and  100°  Centigrade,  is  called  a  calorie.  So  that  250  grams  of 
water  raised  1°  would  represent  250  calories,  and  20  grams  of  water 
raised  5°  would  represent  100  calories. 

While  in  physics  the  unit  of  quantity  is  the  gram,  in  chemistry  the 
unit  which  we  select  is  naturally  that  represented  by  the  formula  of 
the  substance.  Thus,  the  heat  of  combustion  of  carbon  means  the  heat 
produced  by  combining  twelve  grams  of  carbon  with  thirty-two  grams 
of  oxygen,  and  is  sufficient  to  raise  100,000  grams  of  water  one  degree. 
This  is  expressed  as  follows : 

C  +  20  -»  C02  +  100,000  cal. 

In  other  words,  the  combustion  of  less  than  half  an  ounce  of  carbon 
will  raise  one  kilogram  (over  two  pounds)  of  water  from  0°  to  the 
boiling-point 


OXYGEN  77 

It  is  always  found  that  the  same  quantities  of  any  given  chemical 
substances  sustaining  the  same  chemical  change  under  the  same  condi- 
tions produce  or  absorb,  according  as  the  action  is  exothermal  or  en- 
dothermal  (p.  27),  amounts  of  heat  which  are  equal. 

The  rate  at  which  a  given  chemical  action  is  allowed  to  take  place 
has  no  influence  on  the  total  amount  of  heat  consumed  or  produced. 
It  may  not  at  first  sight  appear  obvious  that  rusting  evolves  heat,  but 
a  delicate  thermometer  would  show  that  a  heap  of  rusting  nails  was 
somewhat  higher  in  temperature  than  surrounding  bodies.  Poor 
conductors,  like  oily  rags  and  ill-dried  hay,  show  a  tendency  to  spon- 
taneous combustion  owing  to  accumulation  of  the  slowly  developing 
heat  of  oxidation.  The  warmth  of  our  own  bodies  is  in  part  due  to  the 
same  cause. 

In  accordance  with  invariable  experience  expressed  in  the  law  of 
the  conservation  of  energy,  when  an  action  is  chemically  capable  of 
reversal,  the  contribution  of  the  same  amount  of  heat  which  it  de- 
velops will  exactly  suffice  to  drive  the  chemical  change  in  the  opposite 
direction.  The  heat  contributed  is  simply  used  to  restore  the  amount 
of  chemical  energy  proper  to  the  original  system.  Thus,  the  union  of 
one  chemical  unit  weight  each  of  mercury  and  oxygen  (p.  62)  produces 
30,600  cal. : 

Hg'+  0  ->  HgO  +  30,600  cal., 

and  the  decomposition  of  one  formula-weight  of  mercuric  oxide  (p.  12) 
demands  the  same  amount  of  heat  in  order  that  free  mercury  and  oxy- 
gen, with  their  appropriate  proportions  of  chemical  energy,  may  be 
recovered. 

In  practice  it  is  found  that  all  chemical  changes  are  not  capable  of 
reversal  by  the  use  of  the  sources  of  heat  available  in  the  laboratory. 
A  quantity  of  heat,  equivalent  to  that  produced  by  any  chemical  action 
on  a  small  scale,  is  very  easily  provided,  bat  something  more  appears 
to  be  necessary.  The  heat  provided  must  be  of  a  certain  temperature, 
otherwise  it  is  quite  ineffective.  For  example,  the  heat  produced  by 
the  union  of  calcium  and  oxygen  is  within  the  limits  of  ready  meas- 
urement, 

Ca  +  0  ->  CaO  4-  131,000  cal., 

and  the  supply  of  this  amount  (or  even  of  unlimited  amounts)  of  heat 
to  calcium  oxide  (quicklime)  is  easily  achieved.  Yet  this  method  is 
quite  ineffective  to  produce  decomposition  (p.  63)  of  the  product. 


78  INORGANIC   CHEMISTRY 

Apparently  we  have  not  sufficiently  high  temperatures  for  the  purpose 
at  our  command  (see  Factors  of  energy). 

It  may  be  noted  in  this  connection  that  the  temperatures  required 
to  produce  reasonably  rapid  decomposition  vary  within  a  wide  range. 
Some  substances  can  be  kept  only  below  0°,  and  decompose  when 
allowed  to  become  warm.  Others,  like  the  oxides  of  gold  and  platinum, 
require  a  little  heating  (p.  63).  Many,  like  quicklime,  are  not  broken 
up  even  at  the  temperature  of  the  electric  arc.  When  the  energy  is 
applied  in  the  form  of  electricity  (p.  19),  instead  of  heat,  the  range  is 
incomparably  more  easily  within  the  reach  of  the  means  ordinarily  at 
our  disposal.  There  is  no  substance,  provided  it  is  of  such  a  nature  as 
to  be  affected  by  the  electric  current  at  all,  which  cannot  be  decomposed 
by  a  current  with  an  E.M.F.  of  10  volts  or  less,  while  currents  of  110 
volts  and  over  are  commonly  accessible.  It  is  partly  on  this  account 
that  electrical  processes  have  become  so  common  in  industrial  chem- 
istry. 

One  of  the  most  important  principles  of  thermochemistry  is  the 
law  of  constant  heat  summation.  If  a  system  of  substances  can  be 
transformed  into  another  system  of  substances  by  different  stages  or 
by  more  than  one  route,  then  the  algebraic  sum  of  the  heats  absorbed 
or  produced  in  the  various  stages  is  the  same.  Thus,  barium  oxide 
might  be  formed  directly  from  the  proper  proportions  of  the  constitu- 
ents, or  it  might  be  made  by  preparing  the  dioxide  (p.  63),  and  then 
driving  out  half  of  the  oxygen  contained  in  the  latter.  The  quantities 
of  heat  involved  in  these  changes  are  as  follows : 

Ba  +  O  -»  BaO  +  124,400  cal.  (1) 

(  Ba  4-  2O  ->  Ba02  -f-  141,600  cal.  (2) 

(  BaO,  ->  BaO  +  O  -  17,200  cal.*  (3) 

When  the  last  two  equations  are  added  algebraically,  canceling  terms 
such  as  Ba02  and  O,  which  are  common  to  both  sides  of  the  final 
equation,  the  chemical  action  is  seen  to  be  the  same  as  in  (1),  and  the 
balance,  in  favor  of  heat  produced,  is  124,400  calories  as  before.  If 
in  such  cases  the  sum  of  the  heats  was  not  the  same,  it  would  follow 
that  by  using  different  plans  of  procedure  we  could  prepare  different 
specimens  of  the  same  substances  containing  different  proportions  of 
chemical  energy.  This,  however,  we  have  never  been  able  to  do  (p.  77). 

•  Since  (3)  is  to  be  added  to  (2)  so  as  to  give  (1),  the  formulae  of  (3)  must  be 
so  placed  that  if  the  initial  substances  of  (1)  are  on  the  left  side  of  (2),  the  products 
of  (1)  may  be  on  the  right  side  of  (3)  and  vice  versa. 


OXYGEN  79 

The  quantities  of  heat  liberated  in  two  chemical  changes  are  often 
measures  of  the  relative  amounts  of  available  chemical  energy  in  the 
systems  before  the  change,  and,  therefore,  often  furnish  a  measure  of 
the  relative  chemical  activities  of  the  two  sets  of  substances.  The 
comparison  may  safely  be  made  in  certain  cases  when  the  conditions 
under  which  the  two  actions  take  place  are  precisely  alike.  Formerly 
it  was  supposed  that  the  heat  liberated  was  always  proportional  to  the 
chemical  activity  of  the  substances,  but  we  have  already  shown  cause 
(pp.  27-28)  why  this  general  statement  cannot  be  true. 

Exercises.  —  1.  Construct  the  equations  for  the  combustion  of 
phosphorus,  carbon,  and  iron  in  oxygen  (p.  70). 

2.  When  1  g.  of  sodium  burns  in  oxygen,  it  produces  1.7  g.  of  the 
oxide.     What  is  the  formula  of  the  latter,  and  the  equation  (p.  69)  ? 

3.  Which  are  the  components  (p.  34)  of  the  liquid  made  by  treating 
phosphoric  anhydride  with  water  ?    Which  are  the  constituents  (p.  34) 
of  phosphoric  acid  ? 

4.  How  should  you  show  that,  in  the  making  of  oxygen  from  a 
mixture   of  potassium  chlorate   and  manganese   dioxide,   the   latter 
remains  unchanged  ?     Which  properties  (p.  37)  are  you  employing  for 
this  purpose  ? 

5.  Discuss  the  union  of  iron  and  sulphur  (p.  11)  and  the  decompo- 
sition of  mercuric  oxide  (p.  12)  in  their  relation  to  the  explanations  on 
pp.  73-74. 

6.  How  many  calories  are  required  to  raise  500  g.  of  a  substance 
of  specific  heat  0.5  from  15°  to  37°  (p.  76)? 

7.  The  combustion  of  1  g.  of  sulphur  to  sulphur  dioxide  develops 
2220  calories.     What  is  the  heat  of  combustion  of  sulphur  (p.  76)  ? 

8.  Outline  briefly  the  proof  that  thermochemical  data  are  not  accu- 
rate measures  of  chemical  activity  (p.  79). 


CHAPTEK   VI 
THE   MEASUREMENT    OF   QUANTITY   IN  GASES 

WE  have  spoken  of  measuring  the  proportion  by  weight  of  the 
oxygen  used  in  several  chemical  changes,  but  in  our  illustrations  we 
have  never  weighed  the  gas  itself.  We  have  always  (e.g.  p.  69) 
obtained  its  quantity  by  subtracting  the  weights  of  solid  or  liquid 
bodies.  In  practice  this  method  often  serves  the  purpose. 

Our  preference  for  weighing  as  a  means  of  ascertaining  quantity 
of  matter  is  largely  due  to  the  fact  that  the  weight  is  independent  of 
the  physical  or  even  chemical  condition  of  the  substance.  Yet,  with 
proper  precautions,  we  may  learn  the  quantity  of  matter  by  means  of 
any  other  attributes  which  are  proportional  to  it.  Now  the  volume  is 
such  an  attribute.  In  determining  the  quantity  of  a  liquid,  where 
rapidity  with  no  great  accuracy  is  desired,  the  volume  is  frequently 
measured.  In  the  case  of  gases  the  error  made  in  measuring  the  vol- 
ume is  less,  as  a  rule,  than  in  measuring  the  weight. 

The  Variable  Concentration  of  Gases.  —  A  little  experience 
with  gases  soon  shows  us  that  measurement  of  volume  alone  does 
not  necessarily  give  any  definite  idea  of  the  quantity  of  matter 
present.  The  denseness  with  which  the  gaseous  matter  is  packed  (the 
concentration  of  the  gas)  in  the  vessel  must  somehow  be  defined,  as 
well  as  its  volume,  in  order  that  there  may  be  specification  of  the 
quantity  of  matter. 

Gases  vary  markedly  in  chemical  activity  with  changes  in  their  concentration 
(cf.  p.  75),  and  thus  the  consideration  of  this  condition  (p.  52)  is  continually 
forced  upon  the  chemist.  Solids  and  liquids  do  not  alter  their  denseness  of  pack- 
ing (concentration)  very  noticeably  even  when  compressed  severely  or  changed  in 
temperature.  So  that  concentration  need  not  be  considered  in  the  case  of  pure 
bodies  in  the  solid  or  liquid  forms.  Such  substances  can  be  scattered  through  a 
variable  space  by  solution  in  some  suitable  solvent,  however,  and  then  their  degree 
of  packing  or  concentration  becomes  an  important  factor  in  their  chemical  behavior 
also. 

The  principle  used  for  estimating  the  concentration  of  a  gas  may 
be  illustrated  by  means  of  the  arrangement  in  Fig.  24.  Except  that 


THE   MEASUREMENT   OF   QUANTITY   IN   GASES 


81 


a  little  gas  (any  gas  will  do)  remains  shut  off  by  the  mercury  in  the 
left  limb  of  the  tube,  the  whole  apparatus  has  been  evacuated.  The 
reservoir  can  be  turned  upward,  and  thus  larger  amounts  of  mercury 
may  be  introduced  into  the  tube. 

Now  the  portion  of  the  mercury  below  the  dotted  line  is  essentially 
a  balance,  that  is  to  say,  it  will  move  in  one  direction  or  the  other  if 
the  stresses  on  either  side  change.  At 
present  these  stresses  must  be  equal.  On 
the  right  pan  of  the  balance,  so  to  speak, 
the  stress  is  represented  by  the  weight  of 
the  column  of  mercury  above  the  dotted 
line.  As  there  is  nothing  in  the  tube 
above  this  mercury,  the  weight  of  the 
latter  is  all  that  this  side  of  the  balance 
sustains.  On  the  left  pan  of  the  balance 
there  must  be  an  equal  stress,  and  this 
stress  can  be  caused  only  by  the  gas  con- 
fined in  the  shorter  limb.  The  nature  of 
a  gas  suggests  that  this  stress  must  be 
exercised  on  the  walls  of  the  tube  also, 
although  they  naturally  do  not  exhibit 
its  effects.  This  stress  we  call  the  press- 
ure or  tension  of  the  gas. 

The  height  of  the  surface  of  the  mer- 
cury on  the  right  above  that  on  the  left 
having  been  measured,  more  mercury  may 
now  be  added  from  the  reservoir,  and  the 
difference  in  the  two  levels  again  noted. 
The  gas  cannot  have  diminished  in  amount, 

yet  it  now  occupies  a  smaller  space,  and  is,  therefore,  packed  more 
closely  —  its  concentration  is  greater  than  before.  If,  for  example,  the 
difference  in  level  is  now  twice  as  great,  it  will  be  found  that  the  con- 
centration of  the  gas  is  also  twice  as  great  (its  volume  having  become 
half  of  the  original  volume).  Whatever  amount  of  mercury  is  added, 
we  shall  always  find  that  the  concentration  of  the  gas  is  proportional 
to  the  height  of  the  mercury.  But  this  in  turn  is  proportional  to  the 
weight  of  metal.  The  weight  of  mercury  on  one  side  must,  therefore, 
be  equal  to  the  stress  or  pressure  or  tension  of  the  gas  on  the  other 
side  which  balances  it.  Hence,  the  concentrations  of  a  sample  of  any 
gas  are  proportional  to  the  corresponding  pressures  it  exercises.  We 


FIG.  24. 


82 


INOKGANIC   CHEMISTRY 


determine,  therefore,  the  denseness  with  which  any  sample  of  gas  is 
packed  by  measuring  its  pressure. 

Method  of  Allowing  for  Varying  Concentration  in  Measur- 
ing Quantity  in  Gases.  —  The  principle  just  stated  is  applied  to  the 
measurement  of  the  quantity  of  matter  in  a  sample  of  gas  by  per- 
mitting the  concentration  of  the  sample  to  alter  until  it  is  equal  to 
that  of  the  atmosphere  at  the  moment.     Then  we 
read  off  the  volume   now  occupied,  and  simulta- 
neously we  ascertain   the   pressure  by   observing 
that  of  the  atmosphere.     Each  of  these  two  oper- 
ations is  facilitated  by  a  special  arrangement  of 
apparatus. 

A  gas  to  be  measured  is  always  confined  so 
that  some  liquid  constitutes  one  part  of  the  barrier 
to  its  escape.  The  very  simplest  form  of  the 
apparatus  is  shown  in  Fig.  25.  To  render  the 
concentration  (and  pressure)  of  the  gas  equal  to 
that  of  the  atmosphere,  the  cylinder  containing 
the  gas  is  lowered  (or  raised)  until  the  levels  of  the 
liquid  inside  and  outside  are  the  same.  When  the 
system  is  in  this  condition  the  stress  of  the  gas  on 
the  inner  surface  must  be  equal  to  that  of  the 
atmosphere  on  the  outer  one,  otherwise  movement 
of  the  liquid  would  occur.  The  volume  of  the  gas 
is  then  read  directly  from  the  graduation  on  the 
cylinder.  Often  the  cylinder,  or  other  vessel,  is 
closed  with  a  ground-glass  plate,  placed  quickly  in 
erect  position,  and  weighed.  The  weight  of  water 
which  is  then  required  to  fill  it  to  the  brim  gives  more  exactly  the 
volume  occupied  by  the  gas  (1  g.  water  =  1  c.c.).  When  special 
modes  of  admitting  or  handling  the  gas  have  to  be  provided  for, 
the  apparatus  may  be  more  complex.  But  the  principle  of  adjust- 
ment is  always  the  same.  In  exact  work,  mercury  is  employed  to 
confine  the  gas.  Water  serves  the  purpose  of  rough  work  with 
gases  which,  like  oxygen,  are  but  slightly  soluble  in  it.  When 
water  is  used,  the  volume  is  too  great  by  the  space  occupied  by 
the  vapor  of  the  water  which  is  mixed  with  the  gas  (see  Mixed  gases), 
and  a  correction  must  always  be  made  on  this  account. 

The  pressure  or  tension  of  the  atmosphere  at  any  moment  is  meas- 


PlG.  25. 


THE   MEASUREMENT  OF  QUANTITY  IN   GASES 


83 


ured  by  means  of  a  simplified  form  of  the  apparatus  in  Pig.  24.  The 
reservoir  is  omitted.  The  atmospheric  air  being  the  gas  whose  concen- 
tration is  to  be  measured  through  its  pressure,  the  short  limb  is  left 
open.  The  resulting  apparatus  (Fig.  26)  performs  its  functions  in  the 
same  way  as  does  the  more  complex  one.  The  only  difference  is  that 
mercury  is  automatically  added  to  or  withdrawn  from  the  right  side 
by  the  motion  of  the  metal  resulting  from 
change.s  in  the  pressure  of  the  air.  The 
reading  of  the  vertical  height  between  the 
lower  and  upper  surfaces  of  the  mercury 
gives  a  number  which  is  proportional  to  the 
weight  of  mercury  on  the  right  side  of  the 
balance  and,  therefore,  to  the  (equal)  stress 
of  the  atmosphere  on  the  left.  This  is  called 
the  barometric  reading  (unconnected),  after 
the  name  of  the  instrument. 

To  make  different  readings,  taken  when 
the  mercury  is  at  different  temperatures, 
strictly  proportional  to  the  weight  of  the 
metal,  the  observed  height  is  always  reduced 
to  that  which  would  have  been  shown  by 
the  same  weight  of  mercury  at  0°  in  the 
same  apparatus.  A  thermometer,  and  a  table 
of  temperatures  with  the  corresponding  cor- 
rections to  be  subtracted  from  the  uncorrected 
reading  (0,  Fig.  26),  must  be  used. 

Knowing  now  the  volume  occupied  by 
the  sample  of  gas  when  its  concentration  is 
equal  to  that  of  the  atmosphere  and  the 
barometric  reading,  which  is  proportional  to 
this  concentration,  the  measurement  of  the 
amount  of  matter  in  the  sample  has  become 

definite  so  far  as  concerns  the  variability  of  concentration  with  change 
in  pressure. 

The  recorded  results  of  measurements  made  as  above  at  different 
times  are  still  unsatisfactory  because  the  data  for  samples  of  the  same 
kind  of  gas  differ  in  the  value  of  the  pressure  as  well  as  in  that  of  the 
volume.  To  make  the  results  easily  comparable  in  respect  to  the 
amount  of  matter  they  represent,  one  further  step  is  needed.  All 
the  data  are  recalculated  so  as  to  show  the  volume  each  sample  would 


FIG.  26. 


84  INORGANIC  CHEMISTRY 

have  occupied  if  the  pressure  had  been  equal  to  the  weight  of  760  mm. 
of  mercury,  which  is  the  average  height  of  the  barometer  at  the  sea 
level  in  45°  of  latitude. 

We  have  seen  that  the  concentration  of  a  given  quantity  of  a  gas 
is  proportional  to  its  pressure  (p.  81).  But,  volume  occupied  is  the 
inverse  of  concentration.  Thus  the  same  rule  may  be  stated  in  the 
form,  the  volumes  occupied  by  a  sample  of  any  gas  are  inversely  pro- 
portional to  the  pressure  at  each  volume.  The  fact  was  discovered  by 
Boyle  (1660)  who  stated  it  in  this  way.  In  still  other  words,  th< 
•product  of  the  several  pressures  and  corresponding  volumes  of  a 
sample  of  gas  is  a  constant. 

A  numerical  illustration  will  show  the  mode  of  applying  this  rule. 
We  measure  200  c.c.  of  a  gas  at  atmospheric  pressure,  and  the  barom- 
eter reads  742  mm.  The  question  is :  What  would  be  the  volume 
of  this  amount  of  gas  at  760  mm.  barometric  pressure  ?  It  will  be 
200  x  ^f  $  c.c.  =  volume  at  760  mm.  It  is  unnecessary  to  use  any 
formula,  but  absolutely  essential  to  ask :  Is  the  new  pressure  greater 
or  less  than  the  old  ?  Here  it  is  greater.  Hence,  according  to  the 
law,  the  new  volume  will  be  less,  so  that  the  fraction  must  be  arranged 
with  the  smaller  number  in  the  numerator. 

Boyle's  law  may  be  illustrated  by  using  a  long  tube  bent  like  a  barometer  (Fig. 
26)  but  having  the  short  limb  closed  and  the  long  one  open.  Strips  of  paper  mark 
the  levels  of  the  mercury,  which  are  at  first  alike  on  both  sides,  and  register  the 
volume  of  the  air  in  the  short  limb  at  a  pressure  of  one  atmosphere.  The  reading 
of  the  barometer  at  the  time,  say  750  mm. ,  is  noted.  Then  mercury  is  added  through 
a  funnel  inserted  in  the  long  limb,  until  the  level  in  this  limb  is  750  mm.  above 
that  in  the  other.  A  stick  cut  to  750  mm.  length,  and  held  beside  the  tube,  will 
conveniently  show  when  this  has  been  accomplished.  The  pressure  in  the  open 
limb  being  now  two  atmospheres,  the  volume  of  the  air  will  be  found  to  have 
diminished  to  one-half  its  former  value. 

For  pressures  lower  than  one  atmosphere,  a  different  arrangement  must  be 
used.  A  graduated  tube,  closed  at  one  end,  is  partly  filled  with  mercury  and  in- 
verted in  a  tall,  narrow  cylinder  full  of  the  same  metal.  The  tube  is  then  clamped 
in  any  position,  such  that  the  mercury  level  in  the  tube  is  above  that  in  the  cylinder. 
The  reading  of  the  barometer  is  next  noted.  The  volume  occupied  by  the  air  iu 
the  tube  is  then  read,  and  the  difference  in  height  of  the  two  mercury  surfaces  is 
measured  by  means  of  a  graduated  rule.  Subtracting  this  height  from  that  of  the 
barometer  gives  the  pressure  of  the  air  in  the  tube.  The  position  of  the  tube  is 
then  altered,  and  the  same  measurements  repeated,  as  often  as  is  wished.  The 
product  of  each  volume  and  the  corresponding  pressure  will  be  a  constant  number. 
The  law  is  expressed  mathematically  by  letting  pl  and  v}  represent  the  initial  press- 
ure and  volume,  p2  and  v2  the  new  pressure  and  volume,  and  so  forth.  Then 
plvl  =  p2v2  =  constant  for  that  particular  specimen  of  gas.  For  a  given  sample 
of  gas,  any  one  of  the  four  values  may  be  calculated  if  the  other  three  are  knowr 


THE   MEASUREMENT  OF  QUANTITY  IN   GASES 


85 


Boyle's  law  states  the  facts  with  sufficient  accuracy  for  all  ordinary 
purposes.  But  in  reality  no  two  gases  behave  precisely  alike  in  re- 
spect to  change  in  concentration  when  the  pressure  is  altered  (see  Molar 
weights).  The  same  gas  does  not  even  behave  in  precisely  the  same 
way  at  high,  intermediate,  and  low  pressures.  The 
ideal  gas,  which  should  behave  uniformly,  we 
call  the  perfect  gas.  With  most  gases,  at  low 
pressures  concentration  increases  more,  and  at 
very  high  pressures  much  less  than  the  rule  in- 
dicates (see  Kinetic  hypothesis). 

The  Relation  of  the  Volume  of  a  Gas  to 
Temperature.  —  In  the  foregoing  we  have  as- 
sumed that  there  were  no  temperature  effects. 
But,  as  a  matter  of  fact,  a  rise  in  temperature 
will  at  once  produce  an  increase,  and  a  fall  in 
temperature  a  decrease  in  the  pressure  of  an 
inclosed  sample  of  a  gas.  Hence  a  record  of  the 
pressure  alone  will  fail  to  indicate  the  concentra- 
tion definitely,  and  volume  and  pressure  together 
will  still  leave  the  amount  of  material  unspecified. 
The  temperature  must,  therefore,  be  given  as  well. 

Our  descriptions  of  different  samples  of  gas 
having  thus  become  once  more  incomparable, 
we  apply  the  same  kind  of  remedy  as  before. 
We  calculate  the  volume  which  each  specimen 
of  gas  would  occupy  at  0°. 

The  rule  for  this  calculation  may  be  demonstrated  in  a  rough  way 
as  follows :  The  large,  graduated  bulb  (Fig.  27)  is  surrounded  by  a 
vessel  which  can  subsequently  be  filled  with  ice  water  or  with  water 
of  any  temperature  up  to  100°.  About  one-half  of  its  volume  is  occu- 
pied by  the  gas.  The  mercury  which  fills  the  rest  is  connected  with 
a  reservoir,  so  that  the  levels  of  the  metal  can  be  made  alike,  and  the 
pressure  of  the  gas  be  maintained  constantly  the  same  as  that  of  the 
atmosphere.  When,  now,  the  volume  occupied  by  the  gas  at  0°  is 
read,  and  warmer  water  is  introduced,  we  find  that  the  volume  gains 
7|^  of  its  value  at  0°  for  every  degree  through  which  its  temperature 
rises.  If  it  is  cooled  below  0°,  it  loses  7|7  of  its  volume  at  0°  for 
every  degree  through  which  the  temperature  is  lowered.  Observation 
gives  practically  the  same  value  for  all  gases. 


Fio.  27. 


86 


INORGANIC  CHEMISTRY 


The  following  graphic  method  will  put  these  facts  in  a  clearer 
light.     In  Fig.  28  we  have  on  the  left  a  thermometer  scale  divided 


Temp. 
Cent. 


Vol.  of 
273  c.c. 


Temp. 
Abs. 


-  100° 
99° 
98° 


7° 

6° 

5° 

4° 

3° 

2° 

1° 

0° 

-1° 

-2° 

-3° 

-4° 


-271° 
-272° 
-273° 


373  c.c. 
372  c.c. 
371  c.c. 


373° 
372° 
371° 


275  c.c. 
274  c.c. 
273  c.c. 
272  c.c. 
271  c.o. 


-  275° 

-  274° 

-  273° 

-  272° 

-  271° 


FIG.  28. 


2° 
1° 
0° 


into  degrees  Centigrade.  The  middle  line  represents  the  volumes  of 
a  given  sample  of  gas  which  correspond  with  the  successive  tem- 
peratures. If  we,  for  convenience,  take  a  volume  of  273  c.c.  of  a  gas 
at  0°  and  warm  this  through  1°,  then  at  1°  its  volume,  having  gained 
ff|7  of  its  original  value,  becomes  274  c.c.  At  2°  it  has  gained  another 
y|7  of  the  volume  it  had  at  0°  and  becomes  therefore  275  c.c.,  etc. 
If  cooled  below  0°  it  loses  ^^  of  the  volume,  becoming  272  c.c.,  etc. 

If  the  gas  be  heated  to  100°  it  gains  ££§  of  its  volume  at  0°  and 
becomes  373  c.c.  We  might  infer  that  if  it  were  cooled  through  273° 
from  0°  it  would  lose  f  ^f  of  its  volume,  in  other  words,  it  would  dis- 
appear. This  temperature  has  not  yet  been  reached,  and  in  any  case 
all  gases  would  presumably  liquefy  before  reaching  it.  Our  state- 
ments apply  to  ordinary  temperatures  only,  and  within  them  the  law 
holds  with  considerable  strictness.  Now  the  series  of  numbers  on  the 
middle  line  (Fig.  28)  are  all  273  units  larger  than  the  temperatures 
Centigrade  in  the  line  to  the  left  of  the  figure.  If  we  alter  the  gradua- 
tion of  the  Centigrade  thermometer  by  adding  273  to  each  temperature, 


THE  MEASUREMENT  OF  QUANTITY  IN   GASES  87 

we  secure  the  scale  on  the  right,  which  exhibits  degrees  of  the  very 
same  length  as  before,  but  has  the  additional  advantage  that  the  num- 
bers expressing  temperature  are  the  same  as  those  expressing  volume. 
This  is  what  we  call  the  absolute  scale  of  temperature.  In  this  arti- 
ficial case  we  started  with  273  c.c.  at  0°  C.  (273°  Abs.),  and  the  abso- 
lute temperature  is  here  always  numerically  equal  to  the  volume.  If 
a  different  volume  had  been  taken  at  0°,  then  the  volume  assumed  by 
the  gas  at  each  temperature  would  have  borne  a  constant  ratio  to  the 
volumes  recorded.  Hence,  the  volumes  assumed  by  a  sample  of  gas 
at  different  temperatures,  the  pressure  remaining  constant,  are  in  the 
same  proportion  as  the  corresponding  absolute  temperatures.  This  is 
the  modern  way  of  stating  a  fact  which  was  first  discovered  by  Charles 
of  Paris  (1787).  Obviously,  if  the  volume  remains  constant,  then  the 
pressure  will  be  proportional  to  the  absolute  temperature.  The  bare 
fact  underlying  our  statement  of  the  law  is  that  all  gases  suffer  equal 
increments  (or  decrements)  in  volume  (or  pressure)  for  equal  changes 
in  temperature. 

The  discovery  of  this  fact  is  generally  attributed  to  Dalton,  who  first  published 
an  investigation  of  the  subject,  embodying  this  result,  in  1801,  or  to  Gay-Lussac, 
who  made  a  more  complete  investigation  in  1802.  Inasmuch,  however,  as  in 
Chemistry  we  have  another  important  principle  known  as  Gay-Lussac's  Law,  and 
several  which  are  connected  with  the  name  of  Dalton,  it  is  on  the  whole  fortunate 
that  we  are  justified  in  attributing  this  discovery  to  Charles. 

The  application  of  this  law  may  best  be  illustrated  by  an  example. 
We  obtain  200  c.c.  of  a  gas  at  17°  and  wish  to  know  what  volume  it 
would  occupy  at  0°.  To  answer  the  question,  we  convert  the  Centi- 
grade temperatures  to  the  absolute  scale  by  adding,  algebraically,  273 
to  each.  Thus  200  x  f  £f  =  the  volume  at  0°  required.  No  formula 
is  needed.  We  simply  ask  whether  the  new  temperature  is  higher  or 
lower  than  the  old  one.  Here  it  is  lower.  The  new  volume  will 
therefore  be  smaller  than  the  old  one.  So  we  take  care  to  place  the 
smaller  number  in  the  numerator. 

The  law  may  be  put  in  mathematical  form  thus  :  If  we  make  »0  the  volume  at 
0°  and  «t  the  volume  at  the  temperature  t^  then 


The  number  278  +  ^  is  the  value  of  the  temperature  increased  by  the  number 
273,  and  the  sum  is  spoken  of  as  the  absolute  temperature.  The  intervals  between 
the  degrees  on  this  scale  are  the  same  as  those  on  the  Centigrade  thermometer, 
but  the  numbers  corresponding  to  them  are  all  increased  by  273.  Using  I\  for  the 
absolute  temperature  corresponding  to  tf,  the  last  expression,  above,  becomes  equal 


88  INORGANIC   CHEMISTRY 

to  v0  T£\-s  2\.  Now,  similarly,  for  any  other  temperature,  Tv  there  will  correspond 
another  volume,  v2,  and  the  same  relation  will  hold,  namely,  v2  =  v0  ?fa  T2.  Equat- 
ing both  values  of  v0  and  calling  ^T,  the  coefficient  of  expansion,  a,  we  get, 

JL.  =-  _?L    or   5l  -  ^    or   ±  -  *l 
aT,       aT3  „,       aT3  <>3        T,' 

In  words,  the  volumes  occupied  by  a  sample  of  a  gas  are  proportional  to  tbe  abso- 
lute temperatures.  The  last  formula  enables  us  to  calculate  any  one  of  the  four 
values  when  we  know  the  other  three.  It  is  preferable,  however,  that  beginners 
should  use  the  method  employed  in  the  illustration  given  in  the  preceding  para- 
graph. Simple  mathematical  expressions,  like  the  one  representing  this  law,  are 
not  made  to  save  us  the  trouble  of  remembering  the  law  itself,  and  it  would  be 
unfortunate  if  their  use  led  us  to  forget  it. 

An  instructive  graphic  demonstration  of  the  law  is  given  by  Ost- 
wald  in  his  Principles  of  Inorganic  Chemistry,  p.  75. 

The  behavior  of  gases  in  respect  to  changes  of  temperature  and 
pressure  are  perfectly  independent  of  one  another,  so  that  the  above 
laws  may  be  applied  to  any  example,  either  in  succession,  using  the 
answer  for  the  first  calculation  in  making  the  second,  or  simultane- 
ously. Thus  200  c.c.  of  gas  at  742  mm.  pressure  and  17°  become 
200  X  ft*  x  Hi  =  183-8  c-c-  at  °°  and  76°  mm- 


Mixed  Gases.  —  Two  gases  at  the  same  temperature,  provided 
they  do  not  interact  chemically,  do  not  interfere  with  each  other's 
pressures  when  mixed.  Thus,  if  they  are  forced  into  the  same  volume, 
the  pressure  of  the  mixture  is  equal  to  the  sum  of  those  of  the  com- 
ponents (Dalton's  law,  1807).  The  gases  are  therefore  still  thought 
of  individually,  and  the  share  which  each  gas  has  in  the  total  pressure 
is  called  its  partial  pressure.  This,  like  any  other  gaseous  pressure,  is 
proportional  to  the  concentration  of  the  particular  gas  in  the  mixture. 

For  example,  a  gas  measured  over  water  contains  water  vapor. 
The  partial  pressure  of  this,  called  the  aqueous  tension  (q.v.),  which 
is  definite  for  each  temperature,  must  be  subtracted  from  the  total 
pressure.  The  remainder  is  the  partial  pressure  of  the  gas  being 
measured,  and  this  remainder  is  used  as  the  pressure  of  this  gas  in 
any  calculation. 

Densities  of  Gases.  —  The  application  of  the  laws  of  Boyle  and 
Charles  enables  us  to  express  the  quantities  of  matter  in  samples  of 
gases  in  a  definite  and  readily  comparable  manner.  In  describing 
chemical  changes,  however,  it  is  continually  necessary  to  express 
quantities  of  gases  by  weight.  The  relation  between  volume  and 
weight  for  each  kind  of  gas  must,  therefore,  be  ascertained.  If  we 


THE   MEASUREMENT  OF   QUANTITY   IN   GASES 


know,  for  example,  the  weight  of  one  liter  of  each  gas  at  0°  and  760 
mm.  pressure,  conversion  to  other  weights,  volumes,  temperatures, 
and  pressures  can  be  made.  The  one-thousandth  part  of  this  value, 
the  weight  of  1  c.c.,  is  called  the  density  of  the  gas.  Often,  however, 
the  relative  weights  of  equal  volumes,  with  that  of  air  or  hydrogen  as 
unity,  receive  this  name. 

For  most  chemical  purposes  a  high  degree  of  accuracy  is  not  re- 
quired. Different  arrangements  of  apparatus  are  used  according  to 
circumstances.  The  most  direct  method  is  to  employ  a  light  flask 
provided  with  a  rubber  stopper  and  stopcock  (Fig.  29).  By  means 
of  an  air-pump  the  contents  of  the  flask  are  removed,  and  it  is  weighed. 
This  gives  the  weight  of  the  empty  vessel.  The  gas,  whose  density 
is  to  be  ascertained,  is  then  admitted,  and  care  is  taken  that  it  finally 
fills  the  flask  at  the  pressure  of  the  atmos- 
phere. The  flask  is  closed  and  weighed 
again.  The  increase  represents  the  weight 
of  the  gas.  At  the  same  time  the  tempera- 
ture and  barometric  pressure  are  read. 
The  volume  is  determined  by  displacing 
the  gas  once  more  from  the  flask,  filling 
with  water,  and  weighing  again.  The 
difference  in  weight  between  the  empty 
flask  and  the  flask  full  of  water,  in  grams, 
represents  the  volume  of  the  content  of  the 
flask  in  cubic  centimeters.  This  volume 
is  reduced  to  0°  and  760  mm.  by  the  rules 
discussed  above,  and  we  have  then  a  volume 
of  the  gas  and  the  corresponding  weight. 

To  illustrate,  let  us  suppose  that  the 
volume  of  the  flask  is  200  c.c.  and  that  it 
is  filled  with  oxygen  at  17°  and  742  mm. 
The  weight,  we  will  suppose,  is  found  to  be 
0.26  g.  We  ascertained  (p.  88)  by  calcula- 
tion that  at  0°  and  760  mm.  this  volume 
would  be  183.8  c.c.  The  weight  of  a  liter 
is  given  by  the  proportion  183.8:0.26:: 
1000  :x.  Here  x  =  1.415  g.  When  the 

operation  is  performed  carefully,  and  the  weighing  carried  to  the 
nearest,  milligram  instead  of  the  nearest  centigram,  a  result  more 
Dearly  approaching  the  exact  one  (1.429)  may  easily  be  reached. 


Fio.  29. 


90  INORGANIC  CHEMISTRY 

To  get  the  density  of  oxygen  referred  to  hydrogen  as  unity,  we 
must  divide  the  answer  by  the  weight  of  a  liter  of  hydrogen  (0.08987  g.). 
In  the  above  case  the  quotient  is  15.74.  The  accepted  value  is  15.90. 
The  density  referred  to  air  as  unity  is  similarly  obtained  by  dividing 
by  1.293,  the  weight  of  a  liter  of  air  at  0°  and  760  mm.  pressure. ' 

If  a  suitable  pump  is  not  available,  the  flask,  in  this  case  provided  with  two 
openings,  is  weighed  without  preliminary  exhaustion.  This  gives  the  weight  of 
tfhe  vessel  plus  that  of  the  air  it  contains.  A  continuous  stream  of  the  gas  is  then 
passed  into  the  flask  until  the  air  has  been  completely  displaced.  The  vessel  is 
tken  dosed  and  another  weighing  made.  Finally  the  gas  is  displaced  by  water, 
t*nd  a  third  weighing  taken.  The  temperature  and  barometric  pressure  are  noted 
as  usual.  The  last  weighing  gives  the  volume  as  before.  Knowing  that  one  liter 
of  air  weighs  1.293  g.  at  0°  and  760  mm.,  we  may  calculate  readily  the  weight  of 
the  air  which  the  flask  contained  at  the  observed  temperature  and  pressure.  When 
this  is  subtracted  from  the  number  obtained  in  the  first  weighing,  we  have  the 
weight  of  the  empty  flask.  Subtracting  this  in  turn  from  the  second  weighing,  we 
have  the  weight  of  the  gas.  We  obtain  thus  the  weight  of  a  known  volume  of  the 
gas  at  a  known  temperature  and  pressure,  and  finish  the  calculation  as  before. 

The  values  of  the  densities  of  gases  are  of  great  significance  in  the 
chemical  point  of  view.  A  number  of  them  are  given  in  connection 
with  the  discussion  of  molar  weights  (q.v.). 

Vapor  Densities  of  Liquids  and  Solids.  —  The  densities  of 
vapors  are  as  important  to  the  chemist  as  those  of  gases  and,  solids 
and  liquids  being  more  numerous,  are  even  more  frequently  measured. 

The  apparatus  must  be  specially  adapted  to  the  purpose.  In  its 
simplest  form  (Dumas'  method)  it  consists  of  a  bulb  (160-200  c.c.) 
provided  with  a  long,  narrow  tube  (Fig.  30).  This  corresponds  to  the 
flask  used  for  gases.  The  bulb  is  first  weighed  full  of  air,  and  is  then 
charged  with  a  considerable  amount  of  the  substance,  usually  a  liquid. 
It  is  then  suspended  in  a  bath  whose  temperature  can  be  maintained 
at  some  point  above  the  boiling-point  of  the  substance.  Boiling  water 
(100°)  serves  for  liquids  boiling  below  100°.  When  the  vapor  has  ex- 
pelled all  the  air,  and  is  no  longer  seen  to  issue  from  the  tube,  the 
bulb  contains  nothing  but  the  vapor.  The  tip  of  the  tube  is  then 
sealed  with  a  blow-pipe.  The  temperature  of  the  bath  is  taken  and 
the  barometer  is  read.  When  the  bulb  has  cooled,  it  is  weighed  for 
the  second  time.  The  tip  of  the  tube  is  next  filed  off  under  water, 
which  rushes  in  and  fills  the  bulb  almost  entirely.  A  third  weighing, 
which  should  include  the  portion  of  the  neck  filed  off,  gives,  after 
subtraction  of  the  weight  of  the  vessel,  the  content  of  the  bulb.  From 


THE  MEASUREMENT  OF  QUANTITY  IN  GASES  91 

the  content  and  the  temperature  and  pressure  at  the  time  of  the  first 
weighing,  we  calculate  the  weight  of  air  which  the  bulb  originally 
contained.  When  this  is  subtracted,  we  obtain  the  weight  of  the  empty 
bulb.  Subtracting  that  in  turn  from  the  second 
weighing,  we  get  the  weight  of  the  substance 
which,  in  the  state  of  vapor,  just  filled  the  bulb 
at  the  temperature  of  the  bath  (say  100°)  and 
at  the  pressure  indicated  by  the  barometer. 
From  these  data  we  calculate  the  density  re- 
ferred to  hydrogen  (or  air)  as  unity. 

The  reduction  to  0°  and  760  mm.  pressure 
by  rule  gives,  of  course,  a  fictitious  result. 
The  vapor  would  condense  to  the  liquid  form 
before  0°  was  reached,  if  the  cooling  were  actu- 
ajly  carried  out.  But  the  value  for  the  density 
as  it  ivould  be  at  0°  and  760  mm.  has  to  be  cal- 
culated to  facilitate  comparison  with  the  cor-  FIG.  so. 
responding  values  for  other  substances.  The 

results  have  no  physical  significance,  but  are  highly  important  to  the 
chemist. 

Exercises.  —  The   foregoing  cannot   be  understood  unless  some 
problems  involving  the  laws  of  gases  are  actually  worked. 

1.  Reduce  189  c.c.  of  gas  at  15°  and  750  mm.  to  0°  and  760  mm. 

2.  Eeduce  110  c.c.  of  gas  at  —  5°  and  741  mm.  to  0°  and  760  mm. 

3.  Convert  500  c.c.  of  gas  at  25°  and  700  mm.  to  18°  and  745  mm. 

4.  A  bulb  full  of  air  at  20°  weighs  13.3125  g.     After  being  filled 
with  the  vapor  of  carbon  tetrachloride  at  100°,  it  weighs  13.7969  g. 
Filled  with  water  it  weighs  141.3  g.     The  barometric  reading  (corr.) 
is  755  mm.     What  is  the  vapor   density  referred   to  air  at  0°   and 
760  mm.  ? 

5.  The   density  of  a  substance  referred   to  air  is  3.2.     What   is 
the  density  referred  to  hydrogen  ?     What  will  be  the  volume  occupied 
by  10  g.  of  the  substance  at  20°  and  752  mm.  ? 


CHAPTER  VII 
HYDROGEN 

HYDROGEN,  although,  discovered  by  Paracelsus  in  the  sixteenth 
century,  was  confused  with  other  combustible  gases,  and  its  independ- 
ent nature  was  first  established  by  Cavendish  in  1766.  Somewhat 
later  (1781),  the  latter  showed  that  hydrogen  when  it  burned  gave  water 
vapor,  of  which  he  condensed  a  large  quantity  to  the  liquid  form. 
Taken  in  conjunction  with  Lavoisier's  proof  that  oxygen  was  the 
active  substance  in  the  air  (1777),  this  fact  showed  that  water  was  a 
compound  and  not  a  simple  substance.  The  new  element  was  named 
hydrogen  (Gk.  v8a>p,  water  ;  yeVvav,  to  produce). 


Occurrence.  —  The  free  element  is  found,  mixed  with  varying  pro- 
portions of  other  gases,  in  exhalations  from  volcanoes,  in  pockets  found 
in  certain  layers  of  the  rock-salt  deposits,  and  in  some  meteorites.  The 
air  contains  a  mere  trace  of  it,  not  more  than  one  part  in  30,000.  Its 
lines  are  very  prominent  in  the  spectrum  of  the  sun  and  of  most  stars. 

In  combination,  it  constitutes  about  11  per  cent  of  water.  It  is  an 
essential  constituent  of  all  acids.  It  is  contained  also,  in  combination 
with  carbon,  in  the  components  of  natural  gas,  petroleum,  and  all 
animal  and  vegetable  bodies. 

We  have  seen  (p.  33)  that,  using  the  physical  unit  of  weight,  which  is  the 
same  for  all  substances,  the  element  hydrogen  stands  ninth  in  order  of  plentiful- 
ness.  But  the  chemical  work  that  elements  can  do  should  rather  be  reckoned  by 
the  relative  numbers  of  atomic  weights  which  are  available.  When  Clarke's  num- 
bers are  recalculated  to  this  basis,  and  the  number  of  chemical  unit  weights  of 
oxygen  is  called  100,  hydrogen  assumes  a  position  more  in  harmony  with  its  im- 
portance : 

Oxygen  .     .     100.00  Aluminium.     .     8.57  Iron.     .     .     2.97 

Hydrogen    .       30.10  Magnesium.     .     3.29  Calcium      .     2.81 

Silicon    .     .       28.62  Sodium  .     .     .     3.17  Potassium  .     2.64 

Acids.  —  In  making  hydrogen,  the  acids  are  used  almost  exclu- 
sively. Hence,  some  statements  in  regard  to  their  nature  must  be 
made  before  we  can  use  them  intelligently, 


HYDROGEN  98 

The  common  acids  are  hydrochloric  acid  (HC1,  Aq),  sulphuric  acid 
(H2S04,  Aq),  and  nitric  acid  (HN08,  Aq).  The  usual  forms  are  mix- 
tures containing  water,  the  variable  amount  of  the  latter  being  indi- 
cated by  the  symbol  Aq.  The  first  is  a  solution  of  a  gas,  hydrogen 
chloride.  The  "  pure  concentrated  "  hydrochloric  acid  used  in  lab- 
oratories contains  nearly  as  much  of  the  gas  (39  per  cent  by  weight) 
as  the  water  can  dissolve.  When  heated,  it  readily  gives  up  part  of 
the  gas,  and  the  effervescence  attending  this  must  not  be  mistaken  for 
evidence  of  chemical  action.  The  "  commercial "  acid  contains  im- 
purities and  is  also  less  concentrated.  The  "concentrated"  sulphuric 
acid  is  an  oily  liquid  containing  practically  no  water.  The  "  com- 
mercial "  sulphuric  acid  contains  6  to  7  per  cent  of  water,  besides 
impurities.  The  "  pure  concentrated  "  nitric  acid  contains  70  per  cent 
of  liquid  nitric  acid.  The  "  commercial "  acid,  53  to  62  per  cent  and 
impurities  in  relatively  small  amounts.  Acetic  acid  (HC2H802,  Aq) 
is  a  solution  of  a  liquid  in  water.  All  the  "  dilute  "  acids  contain  90 
to  95  per  cent  of  water.  The  water,  as  a  rule,  takes  no  part  in  the 
chemical  changes  in  which  the  acids  are  concerned,  and  is  therefore 
omitted  from  the  equations. 

The  name  "  acid  "  is  restricted  to  one  class  of  substances  having  cer- 
tain definite  characteristics.  Hydrogen  is  the  only  essential  constitu- 
ent of  all  acids.  When  free  from  water  they  do  not  conduct  electricity. 
Their  aqueous  solutions  have  a  sour  taste  and  change  the  color  of 
litmus  from  blue  to  red.  We  shall  note  presently  two  other  prop- 
erties which  acids  show  when  dissolved  in  water  :  —  They  conduct  and 
are  decomposed  by  the  electric  current,  and  their  hydrogen  (or  one 
unit  weight  of  it  in  the  case  of  acetic  acid)  is  displaced  by  certain 
metals. 

In  describing  the  chemical  behavior  of  acids,  we  speak  of  the 
material  combined  with  the  hydrogen  as  the  negative  radical  (see  next 
section).  Thus  the  negative  radicals  in  the  above  acids  are  Cl,  S04, 
N08,  and  C2H302,  respectively.  The  first  (Cl)  is  a  simple  radical,  the 
others  complex.  In  many  interactions  the  complex  radicals  move  as 
units  from  one  state  of  combination  to  another. 

Preparation  of  Hydrogen  by  Electrolysis.  —  A  supply  of 
free,  unmixed  hydrogen  not  existing,  we  are  compelled  to  prepare  it 
from  compounds  by  the  two  general  plans  (p.  63)  used  for  liberating 
simple  substances.  When  the  first  plan,  the  direct  application  of 
energy,  is  employed,  we  find  that  electricity  serves  the  purpose  best. 


94 


INORGANIC  CHEMISTRY 


The  common  compounds  of  hydrogen,  like  hydrogen  chloride  and  water,  are 
not  easily  decomposed  by  heat,  and  in  most  cases,  at  best,  a  mixture  of  gases  would 
be  obtained.  The  difficulty  in  separating  the  resulting  gases  makes  the  use  of  this 
form  of  energy  unsuitable.  On  account  of  its  ability,  not  only  to  liberate  the  con- 
stituents from  combination,  but  also  to  deliver  the  positive  and  the  negative  parts 
of  the  compound  in  separate  places,  electricity  alone  is  available. 
to 

If  we  dissolve  any  acid  in  water,  and  immerse  the  wires  from  a 
battery  in  the  solution,  bubbles  of  hydrogen  begin  to  appear  on  the 
negative  wire  (the  cathode)  and  rise  to  the  surface.  All  the  other  con- 
stituents, whatever  they  may  be,  are 
attracted  to  the  positive  wire  (the  anode) 
and  are  set  free  in  some  form  at  its 
surface.  It  is  on  account  of  this  be- 
havior of  the  radicals  of  acids  that  they 
are  known  as  "  negative  "  radicals.  An 
apparatus  devised  by  Hofmann  (Fig.  31) 
enables  us  to  secure  the  hydrogen,  which 
ascends  on  the  left  and  accumulates  at 
the  top  of  the  tube,  displacing  the  solu- 
tion. The  other  products,  if  gaseous, 
occupy  a  separate  tube  on  the  right 
side.  The  solution  displaced  by  the 
gases  is  forced  down  and  mounts  into 
the  bulb  behind.  The  current  of  elec- 
tricity flows  from  one  wire  to  the  other 
through  the  cross-tube.  In  the  typical 
case,  with  a  properly  selected  acid,  the 

O/         ir^  production   of   hydrogen    ceases    when 

$9^  ^^  the  acid  is  all  decomposed.     The  water 

'  alone   is   an   almost   complete  noncon- 

ductor, so  that  the  flow  of  the  electricity 
practically  ceases  at  the  same  time.     If 
FIG.  si.  the  operation  does  not  come  to  rest  in 

this  way,  its  continuance  is  due  to  the 

regeneration  of   conducting   substances   by  the  interaction  with  the 
water  of  the  materials  of  the  radical  liberated  at  the  positive  electrode. 
When  hydrochloric  acid  is  used,  we  have  a  close  approximation  to 
the  typical  case.     The  equation  is  : 

HC1  — >  H  (neg.  wire)  +  Cl  (pos.  wire), 
and  the  chlorine,  a  soluble  gas,  remains  dissolved  in  the  water  near  one 


HYDROGEN  96 

pole.     When  sulphuric  acid  is  employed,  the  equation  is : 

H2S04  -»  2H  (neg.  wire)  +  S04  (pos.  wire), 

and  the  S04  interacts  with  the  water  (see  Discharging  potentials),  thus  : 
S04  H-  H20  -i  H2S04  +  0. 

Hence  oxygen  comes  off,  and  the  substance  regenerated  is  here  sul- 
phuric acid  itself.  The  final  results  are,  therefore,  the  liberation  of 
hydrogen  and  of  oxygen  and  the  localization  of  the  regenerated  acid 
round  the  positive  electrode. 

It  is  worth  noting  that  the  acids  and  water,  taken  separately,  are 
all  nonconductors.  The  fact  that  the  mixture  does  conduct,  concomi- 
tantly  with  the  decomposition  of  the  acid,  is  therefore  highly  sug- 
gestive. Solution  in  such  cases  must  be  something  more  than  a  mere 
physical  change  of  state  of  aggregation  (see  lonization). 

It  is  commonly  asserted  that  water  is  decomposed  by  a  current  of  electricity 
This  is  true  in  the  sense  in  which  we  might  say  that  a  man  can  carry  off  a  hill. 
He  may  eventually  remove  it,  if  you  give  him  time.  The  action  of  electricity  upon 
the  purest  water  is  exceedingly  slow,  on  account  of  the  very  minute  conductivity 
for  electricity  which  it  possesses.  Common  distilled  water  owes  its  appreciable 
capacity  for  conducting  chiefly  to  traces  of  an  acid,  namely,  carbonic  acid,  which 
it  contains.  Even  when  the  water  is  saturated  with  carbonic  acid,  however,  dilute 
sulphuric  acid  has  a  conductivity  of  the  order  of  a  thousand  times  greater.  For 
our  present  purpose,  therefore,  water  is  declared  to  be  a  nonconductor.  Yet,  as 
we  shall  see,  the  conductivity  of  pure  water,  small  as  it  is,  has  to  be  taken  into 
consideration  in  certain  cases  (see  Hydrolysis  and  Electromotive  chemistry). 

Preparation  of  Hydrogen  by  Displacement  from  Diluted 
Acids.  —  By  use  of  the  second  plan  for  liberating  elements  (p.  63),  hy- 
drogen may  be  obtained  from  acids,  through  substitution  for  it  of  some 
element  with  which  the  negative  radical  will  unite. 

The  acids  must  be  diluted  with  water  before  rapid  action  occurs. 
The  substances  capable  of  displacing  hydrogen  from  them  are  certain 
of  the  metals,  like  zinc,  iron,  and  aluminium.  The  hydrogen  escapes 
in  bubbles,  and  evaporation  of  the  remaining  liquid  gives  in  dry  form 
the  compound  of  the  metal  with  the  other  constituents  of  the  acid. 
Thus,  with  zinc  and  sulphuric  acid,  zinc  sulphate  is  produced : 

Zn  +  H2S04  ->  2H  +  ZnS04 ; 

and  with  tin  or  aluminium  and  hydrochloric  acid  we  get  stannous 
chloride  or  aluminium  chloride : 


96 


INORGANIC   CHEMISTRY 
Sn  +  2HC1  ->  2H  +  SnCl2 

(Tin)  (Stann< 


(Tin) 


(j.m; 

Al  +  3HC1  -»  3H  +  A1C1, 


9^JLA.\^J.n 

(Stannous  chloride) 


The  wate»  undergoes  no  change  during  the  action,  although  its  presence 
is  essential.  It  is  simply  a  part  of  the  apparatus.  Any  acid  may  be 
used,  although  with  many  the  action  goes  on  very  slowly.  In  all 

cases  the  plan  of  the  action  is  the 
same  :  the  metal  is  said  to  displace 
the  hydrogen  (see  below). 

The  apparatus  for  generating 
small  amounts  of  hydrogen  (Fig.  32) 
is  arranged  so  that  additional  acid 
may  be  added  through  the  thistle 
or  safety  tube.  This  avoids  ad- 
mission of  air.  With  a  Kipp's 
apparatus  (Fig.  33)  the  gas  may  be 
made  on  a  larger  scale  and  its  de- 
livery can  be  regulated.  When  the 
stream  of  gas  is  shut  off  by  the 
stopcock,  the  pressure  of  the  gas, 
as  it  continues  to  be  generated, 
drives  the  acid  away  from  the  metal 
and  up  into  the  globe  above,  so 
that  the  action  ceases.  Yet  the 
action  is  ready  to  begin  again  the 
moment  any  portion  of  the  stored 
gas  is  drawn  off  for  use. 

A  rather  sharp  line  can  be  drawn  between  those  metals  which  dis- 
place hydrogen  from  dilute  acids  and  those  which,  like  mercury, 
silver,  and  gold,  do  not  (see  Electromotive  series  of  the  metals). 

Contact  of  the  zinc  or  iron  with  an  inactive  metal,  like  platinum, 
always  hastens  the  interaction,  and  therefore  renders  the  evolution  of 
the  hydrogen  more  conspicuous.  Such  an  arrangement  is  called  a 
couple,  and  its  efficiency  depends  on  the  electric  states  of  the  two 
metals  (see  Solution  tension). 

When  water  is  not  used  along  with  the  acid,  the  latter  is  either 
inactive  or  undergoes  a  different  sort  of  chemical  change.  Thus,  dry, 
gaseous  or  liquefied  hydrogen  chloride  hardly  interacts  at  all  with 
zinc.  Pure,  concentrated  sulphuric  acid,  on  the  other  hand,  although 
almost  unaffected  by  zinc  in  the  cold,  is  violently  decomposed  when 


FIG.  32. 


HYDROGEN 


97 


heated.  The  action,  however,  is  not  a  simple  displacement  (see  below) 
of  the  hydrogen.  The  oxygen  is  removed  from  a  part  of  the  acid,  and 
water  and  hydrogen  sulphide  are  formed  : 


4Zn 


5H2S04 


4ZnS04 


4H20 


H2S. 


Preparation  of  Hydrogen  from  Water.  —  Every  one  of  the 
metals  which  act  on  dilute  acids  will  also  displace  hydrogen  from 
water,  and  no  others  will  do  so.  Only 
the  more  active  metals,  like  potassium 
and  sodium,  which  would  act  with 
uncontrollable  vigor  on  dilute  acids, 
can  displace  the  hydrogen  rapidly 
from  cold  water.  Magnesium  and  zinc 
show  obvious  action  on  water  at  100° 
only,  and  are  much  assisted  by  contact 
with  another  metal.  If  any  action  is 
to  be  perceived  in  the  cold,  the  iron, 
nickel,  zinc,  and  magnesium  have  to  be 
used  in  a  state  of  fine  powder,  with 
great  surface. 

In  all  cases  in  which  cold  or  boil- 
ing water  is  employed,  the  hydrogen 
of  the  water  is  not  completely  dis- 
placed. The  metal  forms  an  hydrox- 
ide, such  as  sodium  hydroxide  or 
magnesium  hydroxide  : 

Na  +  H90  ->  H  +  NaOH, 
Mg  +  2H20  -*  2H  +  Mg(OH)2. 


FIG.  33. 


Sodium,  which  is  one  of  the  con- 
stituents of  common  salt,  may  be  used 
to  illustrate  this  sort  of  action.  As 
it  is  lighter  than  water,  it  must  be 

held  under  the  surface  by  means  of  a  piece  of  wire-gauze  in  order 
that  the  gas  may  be  collected  (Fig.  34).  Most  of  the  water  serves 
the  mechanical  purpose  of  permitting  the  collection  of  the  gas,  and 
only  a  small  fraction  of  it  takes  part  in  the  change.  The  solution 
has  a  soapy  feeling  and  turns  litmus  from  red  to  blue.  This  color 
icaction  is  the  precise  opposite  of  that  of  acids  (p.  93).  Substances 


98 


INORGANIC  CHEMISTRY 


causing  these  two  effects  are  called  alkalies.  Evaporation  of  the 
resulting  very  dilute  solution  reveals  the  sodium  hydroxide,  the 
alkali,  as  a  white  solid. 

With  steam  at  a  red  heat,  metals  like  iron,  zinc,  and  magnesium 


FIG.  34. 


interact  vigorously.     The  metal  is  placed  in  a  tube  in  which  it  can  be 
strongly  heated  (Fig.  35X     The  steam,  generated  in  a  flask,  enters  at 


FIG.  35. 


one  end  of  the  tube,  and  the  hydrogen  passes  off  at  the  other.     Since, 
at  a  red  heat,  all  hydroxides,  except  those  of  potassium  and  sodium, 


HYDROGEN  99 

are  decomposed  into  an  oxide  of  the  metal  and  water,  Mg(OH)2  —  > 
MgO  +  H20,  the  oxides  are  formed  in  this  case  : 


Iron  gives  the  magnetic  oxide,  Fe804.  Hence,  to  make  the  equation, 
four  unit-  weights  of  oxygen,  and  therefore  four  formula-  weights  of 
water,  are  required  : 

4H2O  +  3Fe  ->  Fe3O4  +  8H. 

Tests.  —  The  effect  of  a  solution  of  an  acid  (p.  93)  or  an  alkali  (see 
above)  upon  litmus  solution  is  an  illustration  of  a  chemical  teat.  In 
this  case  a  mere  trace  of  a  highly  colored  body  dissolved  in  water 
interacts  with  a  correspondingly  minute  trace  of  the  acid  or  alkali 
under  examination.  Yet  the  product  has  such  marked  color  that  its 
presence  is  immediately  recognizable.  Tests  for  various  kinds  of 
materials  are  eagerly  sought  by  chemists  for  use  in  identifying  un- 
known substances.  Any  property  which  is  so  conspicuous  as  to  be 
apparent  when  very  little  material  is  concerned  will  serve  as  a  test. 
Thus  the  precipitation  of  calcium  carbonate  (p.  68)  enables  us  to 
recognize  the  presence  of  carbonic  acid.  The  test  consisted  in  the 
addition  of  lime-water. 

The  Other  Ways  of  Preparing  Hydrogen.  —  For  special  pur- 
poses, hydrogen  may  be  made  by  boiling  an  aqueous  solution  of  sodi- 
um hydroxide  with  aluminium  turnings,  when  sodium  aluminate  is 
formed  :  Al  +  NaOH  +  H20  ->  NaA102  +  3H  ;  also  by  heating  pow- 
dered zinc  and  dry  sodium  hydroxide,  the  product  being  sodium  zin- 
cate  :  Zn  +  2NaOH  ->  Na2Zn02  +  2H.  It  may  likewise  be  made  by 
electrolyzing  solutions  of  compounds  of  metals  which,  when  free,  dis- 
place hydrogen  from  cold  water.  Thus  electrolysis  of  sodium  chloride 
in  aqueous  solution  liberates  chlorine  at  the  positive  wire  and  hydrogen 
and  sodium  hydroxide  (p.  97)  at  the  negative  wire  (see  ^Discharging 
potentials,  Chap,  xxxviii). 

Displacement.  —  We  now  have  before  us  illustrations  of  two  sub- 
varieties  of  the  third  kind  (p.  15)  of  chemical  change.  In  this  kind, 
compounds  are  decomposed  and  the  parts  combine  in  a  new  way. 
The  first  sub-variety  was  double  decomposition,  as  in  the  action  of 
sodium  chloride  upon  silver  nitrate  : 

NaCl  +  AgNO,  -+  AgCl  + 


100 


INOKGANIC   CHEMISTRY 


In  this  class  of  cases  two  compounds  interact,  each  splits  into  the 
radicals  of  which  it  is  composed,  and  two  new  compounds  are  formed 
by  union  of  the  radicals  crosswise.  The  actions  used  in  the  prepara- 
tion of  hydrogen  differ  from  these  inasmuch  as  one  compound  and  one 
element  interact,  the  compound  splits  into  its  radicals,  and  one  com- 
pound and  one  free  element  are  produced: 

Zn  +  H2S04->  ZnSO,  +  2H, 
Zn  +  2NaOH  ->  Na2Zn02  -f  2H. 

The  former  element,  here  the  zinc,  is  said  to  displace  the  latter,  here 
the  hydrogen,  from  combination.  In  double  decomposition  there 

is  an  even  exchange,  the  sodium, 
for  example,  giving  up  one 
radical  (Cl),  and  getting  another 
(N03),  whereas  in  displacement 
one  element  gains  a  radical 
while  another  loses  it,  the  zinc, 
for  example,  giving  up  nothing 
but  getting  S04,  while  the 
hydrogen  loses  SO4,  and  gains 
nothing  in  return. 

It  should  be  noted,  that 
although  in  the  former  of  the 
above  illustrations  the  Ag  may 
be  said  to  displace  the  Na  from 
combination  with  the  Cl,  the 


FIG.  36. 


displaced   escapes   in   the   free 


term   displace   is 
cally  only  when 
condition. 


used  techni- 
the   element 


Purification  of  Gases.  —  Hydrogen  made  in  any  of  the  above 
ways  is  impure  (p.  34).  As  made  by  the  first  three  methods,  a  good 
deal  of  water  vapor  is  mixed  with  it.  Other  impurities,  like  hydrogen 
sulphide  and  arsine,  come  from  the  action  of  the  acid  on  foreign  mate- 
rials in  the  zinc  (p.  95).  Some  of  the  acid,  if  it  is  volatile,  will  also  be 
taken  over  with  the  gas.  When  the  object  for  which  the  gas  is  being 
made  demands  it,  we  must  know  what  the  impurities  to  be  expected 
are,  and  take  proper  means  of  removing  them. 

Gases  are  freed  from  aqueous  vapor  by  means  of  calcium  chloride 
or  concentrated  sulphuric  acid,  which  greedily  absorb  moisture.  The 


HYDROGEN 


101 


former  is  used  in  granulated  form  in  straight  or  bent  tubes  (Fig.  36). 
The  latter  is  applied  by  saturating  pieces  of  pumice-stone  with  the 
acid  and  filling  similar  tubes  with  the  fragments.  Or  the  acid  may 
be  placed  in  a  gas  washing  bottle  (Fig.  37).  For  extremely  complete 
drying,  a  tube  may  be  filled  with  phos- 
phoric anhydride  sifted  upon  glass 
beads  or  glass  wool.  Forethought  must 
be  used  to  avoid  a  drying  agent  which 
will  interact  with  gas.  The  longer  the 
gas  remains  in  contact  with  the  drying 
agent,  the  more  perfect,  up  to  a  certain 
limit,  is  the  purification  effected.  In 
all  cases,  the  stream  of  gas  must  pass 
slowly. 

Particles  of  liquid  or  solid  matter 
are  always  carried  along  by  freshly 
made  gases.  These  will  pass  with  the 
gas  through  sulphuric  acid  without 
being  affected.  A  plug  of  cotton  or 
of  glass  wool  in  some  part  of  the  tubing 
is  required  to  arrest  them. 


FIG.  37. 


Valence.  —  We  shall  gain  much 
help  in  the  making  of  equations  if  we 
now  introduce  and  bring  into  relation  to 

the  symbols  a  conception  for  which  the  remarks  about  atomic  weights 
(p.  50)  have  paved  the  way.  It  will  have  been  observed  that  the  com- 
position of  the  chlorides  of  aluminium,  tin,  and  sodium  are  represented 
by  the  formulae  A1C13,  SnCl2,  and  NaCl  respectively.  Again,  the  hy- 
droxide of  sodium  is  NaOH,  while  those  of  magnesium  and  calcium  are 
Mg(OH)2  and  Ca(OH)2 .  In  making  equations  we  constantly  need  to 
know  whether  the  chloride  of  an  element,  say  magnesium,  is  MgCl,  or 
MgCl2,  or  MgCl3,  or  MgCl4,  etc.,  and  whether  its  sulphate  is  MgS04, 
or  Mg2S04,  or  some  other  combination  of  the  symbols.  To  answer 
questions  like  this  it  is  not  necessary  to  know  the  formula  of  every 
compound  of  each  element :  the  apparent  disorder  of  these  numbers 
can  be  reduced  to  rule,  and  the  reader  should  endeavor  thoroughly  to 
master  the  rule  before  going  farther. 

If  the  method  by  which  the  atomic  weights  were  derived   from 
equivalents  (p.  50)  is  now  reexamined,  the  nature  of  this  rule  will  be 


102  INORGANIC   CHEMISTRY 

seen.  It  was  found,  for  example,  that  9.03  parts  of  aluminium  (p.  50) 
combined  with  the  equivalent  weights  of  the  other  elements,  and 
therefore  with  35.45  parts  of  chlorine.  If  this  weight  of  aluminium 
had  been  accepted  as  the  final  unit  (the  atomic  weight),  then  it  would 
have  been  represented  by  the  symbol  Al,  and,  since  01  stands  for 
35.45  parts  of  chlorine,  the  formula  of  the  chloride  would  have  been 
A1C1.  In  point  of  fact,  however,  a  number  three  times  as  large  as  the 
equivalent,  namely,  27.1,  was  chosen  as  the  atomic  weight  of  alumin- 
ium, and  symbol  Al  stands  for  this  triple  quantity.  If  the  equivalent 
of  chlorine  had  also  been  tripled  in  making  its  atomic  weight,  the 
amounts  represented  by  the  symbols  would  still  have  been  chemically 
equivalent,  and  the  formula  would  still  have  been  A1C1.  But  the 
equivalent  of  chlorine  was  left  unaltered.  Hence,  to  get  the  equiva- 
lent amounts  (i.e.,  the  actual  combining  quantities)  of  the  two  elements, 
we  must  have  3C1  with  1A1.  The  formula  is  thus  A1C18,  Now,  it 
is  evident  that  this  tripling  of  the  equivalent  of  aluminium  will  affect 
the  formulae  of  all  its  compounds.  Whenever  it  is  combined  with  an 
element  which,  like  chlorine,  has  identical  equivalent  and  atomic 
weights,  the  formula  of  the  compound  will  be  of  the  form  A1X8.  In 
accordance  with  this  we  have  the  bromide  AlBr3.  In  making  the 
formulae  of  compounds  of  aluminium,  the  chief  thing  to  be  kept  in 
mind,  therefore,  is  the  fact  that  its  atomic  weight  contains  three 
equivalents  and  always  combines  with  three  equivalents  of  another 
element.  This  fact  we  state  by  saying  that  the  valence  of  the  atomic 
weight  of  aluminium  is  three,  or  simply  that  the  element  aluminium  is 
trivalent. 

Similarly,  the  equivalent  of  tin  is  59.5  and  its  atomic  weight  is 
119.  This  atomic  weight  therefore  contains  two  equivalents  of  tin 
and  combines  with  two  equivalents  of  any  other  element.  Hence, 
the  formula  of  a  compound  of  tin  with  an  element  of  the  chlorine 
class  will  be  SnX3.  Thus  tin  is  bivalent.  In  like  manner  the  equiva- 
lent of  sodium  is  23,  and  this  number  was  not  altered  in  making  the 
atomic  weight.  Hence,  the  symbol  Na  stands  for  one  equivalent,  and 
the  formula  of  the  compound  with  chlorine  is  NaCl.  Elements  whose 
atomic  weights  are  identical  with  their  equivalents  are  described  as 
univalent. 

Thus  the  valence  of  an  element  may  be  defined  as  the  number  of 
equivalent  -weights  contained  in  its  atomic  weight.  Arithmetically 
it  is  the  integer  by  which  the  equivalent  weight  was  multiplied  in 
forming  the  atomic  weight,  The  above  explanation  shows  that 


HYDROGEN  103 

we  may  define  the  valence  of  an  element  also  as  the  number  of 
atomic  weights  of  a  univalent  element,  with  which  its  atomic 
weight  will  combine.  A  more  complete  definition  will  be  given  pres- 
ently. 

Sometimes  the  valence  is  indicated  in  the  symbol  thus:  A1HI> 
Sn11,  Na1,  Cl1,  Br1.  The  table  of  atomic  weights  (p.  50)  shows  the 
following  additional  cases:  0",  Cu11,  Sn,  Hg11,  H1,  Fe"  Mgn,  CIV. 
With  the  help  of  this  list  the  formulae  of  compounds  may  easily  be 
made.  Thus,  oxygen  is  bivalent,  and  an  atomic  weight  of  oxygen, 
represented  by  0,  will  combine  with  two  atomic  weights  of  a  univalent 
element  as  in  0IIH2I  (water),  or  with  one  atomic  weight  of  a  bivalent 
element  as  in  0"Snn  (stannous  oxide),  0IIHg"  (mercuric  oxide),  OnCun 
(cupric  oxide),  0IIMg11  (magnesium  oxide).  Again,  carbon  being  quad- 
rivalent, the  atomic  weight  combines  with  four  units  of  chlorine  and 
of  hydrogen  in  C^Cl/  (carbon  tetrachloride)  and  CIVH4I  (methane), 
or  with  two  units  of  oxygen  and  of  sulphur  in  CIV02n  (carbon  dioxide) 
and  CIVS2U  (carbon  disulphide).  When  it  combines  with  a  trivalent 
element,  equal  numbers  of  equivalents  of  each  element  must  be  used, 
as  in  C8IVAl4m  (aluminium  carbide),  where  C8  and  A14  contain  twelve 
equivalents  each.  This  method,  with  exceptions  to  be  noted  below, 
will  give  the  formulas  of  all  compounds  containing  only  two  elements 
—  so-called  binary  compounds. 

The  above  mode  of  handling  valence  is  based  upon  the  notion  of 
combination  in  equivalent  proportions.  Another  variety  of  chemical 
change,  namely  displacement  (p.  99),  is  often  of  assistance  in  enab- 
ling us  to  determine  the  valence  of  an  element.  It  will  be  noted 
that  when  Al  acted  upon  hydrochloric  acid  (p.  96)  and  combined 
with  301,  it  necessarily  displaced  the  3H  with  which  the  301  was 
formerly  united.  It  was  equivalent  to  3H  for  the  purpose  of  holding 
301  in  combination.  It  is  from  this  aspect  of  the  relation  that  the 
word  "  valence  "  comes.  Al  is  equi-valent  to  3H,  and,  H  having  the 
unit  valence,  Al  is  trivalent.  Similarly,  since  one  atomic  weight  of 
zinc,  represented  by  the  symbol  Zn,  displaces  2H  (p.  95),  zinc  must 
be  bivalent.  Combining  this  with  the  former  conception,  we  reach  the 
common  definition  of  the  valence  of  an  element :  The  valence  of  the 
atomic  'weight  of  an  element  is  the  number  of  atomic  weights  of  hydro- 
gen, or  of  some  other  univalent  element,  which  it  combines  with  or 
displaces. 

Formulae  are  often  written  so  as  to  show  the  valence  plainly. 
Thus,  K — 01  indicates,  by  the  single  line,  that  each  element  is  uni- 


104  INORGANIC   CHEMISTRY 

valent.  Two  or  more  lines  meeting  at  a  symbol,  indicates  that  that 
element  is  bivalent  or  tri  valent : 

/  Cl  p1  Fp=O 

A1-C1,    Mg^{,    Mg  =  0,    0  =  Fe-0-Fe=0,        >Q. 

\C1          x°  Fe=0 

Here  Al  and  Fe  are  trivalent,  Mg  and  O  bivalent.  The  lines  may  be 
drawn  in  any  direction,  as  the  last  two  formulae  show.  As  many 
lines  proceed  from  each  symbol  as  will  represent  its  valence.  The 
resulting  structures  are  called  graphic  formulae. 

The  Valence  of  Radicals.  —  In  the  preceding  section  it  has  been 
seen  that  the  valence  of  elements  can  easily  be  determined  when  they 
are  present  in  binary  combination.  This  is  no  longer  the  case  when 
more  than  two  elements  are  united  together.  A  study  of  chemical 
changes  shows,  however,  that  even  here  the  conception  of  valence  can 
still  be  employed.  In  the  interaction  of  zinc  with  dilute  sulphuric 
acid: 

Zn  +  H,S04-+ZnS04-f  2H 

the  group  S04  passes  as  a  whole  from  combination  with  2H  to  com- 
bination with  Zn.  Hence,  although  we  cannot  by  inspection  deter- 
mine the  valence  of  sulphur,  we  do  perceive  that  the  radical  SO^,  as  a 
whole,  must  be  bivalent.  It  occurs,  in  fact,  in  all  sulphates,  as  Ag2S04, 
MgSO^,  and  A12(S04)3,  and  in  the  interactions  of  these  substances  it 
usually  passes  intact  from  one  state  of  combination  to  another,  and 
behaves  as  if  it  were  a  unit  of  a  single  element  of  valence  two.  Again, 
in  the  interaction  of  salt  with  silver  nitrate  (p.  99),  we  observe  that 
the  radical  NO8  is  univalent.  Still  again,  the  compositions  of  the  com- 
pounds CaCl2  and  Ca(OH)2  show  that  the  radical  OH  (hydroxyl)  is 
univalent.  The  formula  NaOH  leads  to  the  same  conclusion. 

This  addition  to  our  ideas  enables  us  greatly  to  extend  the  list  of 
substances  of  which  we  can  write  the  formulae.  Thus,  the  hydroxides 
all  contain  (OH)1,  e.g.  A1III(OH)8I  (aluminium  hydroxide),  Snn(OH)2I 
(stannous  hydroxide),  Cu^OH),1  (cupric  hydroxide).  The  nitrates 
all  contain  (NO,,)1,  as  :  H^NOg)1  (nitric  acid),  Mgn(N08)2I  (magnesium 
nitrate).  It  is  to  preserve  the  identity  of  the  radicals  that  we  write 
them  in  brackets  and  place  the  factor  outside,  instead  of  using  the 
forms  A103H8,  MgN,O6,  and  so  forth.  In  fact,  we  regard  as  binary 
compounds,  substances  which  commonly  interact  as  if  the  radicals 
were  single  elements.  In  all  actions  in  which  the  radicals  preserve 


HYDROGEN  105 

their  integrity,  the  conception  of  valence  proper  to  binary  compounds 
may  be  used  for  the  more  complex  compounds  also. 

How  to  Ascertain  the  Valence  of  an  Element  or  Radical.  — 

The  above  shows  that  the  valence  of  one  element  or  radical  may  al- 
ways be  ascertained  by  examination  of  the  formula  of  a  compound 
containing  another  element  or  radical  of  known  valence.  Thus,  when 
we  know  the  formula  of  sodium  iodide  to  be  Na1!,  or  that  of  hydro- 
gen iodide  to  be  H1!,  we  infer  that  iodine  is  univalent.  The  formula 
of  silica  (sand)  Si02n  shows  silicon  to  be  quadrivalent,  and  indicates 
that  the  chloride  must  be  SiCl4.  Similarly  the  formula  of  calcium 
carbonate  CanC08  shows  that  the  radical  C03,  which  is  common  to  all 
carbonates,  must  be  bivalent.  Hence,  the  chemist  does  not  memorize 
the  valences  themselves ;  he  recovers  them  when  needed  by  recalling 
the  formula  of  a  compound  containing  a  more  familiar  element  or 
radical. 

It  is  absolutely  essential  that  correct  valences  should  be  used 
in  constructing  equations,  and,  at  first,  the  student  will  find  the  task 
by  no  means  easy.  He  should  give  special  attention  to  this  mat- 
ter until,  by  solving  the  exercises  at  the  end  of  this  chapter,  and  by 
careful  examination  of  all  the  equations  encountered  in  the  text,  he 
has  mastered  the  subject. 

Multiple  Valence  and  Exceptional  Cases.  —  Some  elements 
show  more  than  one  valence.  This  is  as  much  as  to  say  that  an  atomic 
weight  of  such  an  element  may  form  stable  compounds  with  two,  or 
even  more  different  numbers  of  equivalents  of  another  element.  This 
fact  has  already  been  mentioned,  for  it  is  implied  in  the  law  of  mul- 
tiple proportions  (p.  41).  Thus  an  atomic  weight  of  tin  may  form 
two  different  compounds  with  chlorine,  namely,  Sn"Cl2  (stannows 
chloride)  and  SnIVCl4  (stannic  chloride).  Tin  behaves  in  the  same  way 
towards  other  elements,  however,  and  we  have  a  series  of  stannows 
compounds,  SnO,  SnBr2,  and  so  forth,  and  a  corresponding  series  of 
stannic  compounds,  SnO2,  SnBr4,  etc.  Two  different  valences  of  the 
same  element  or  radical  gives  rise  therefore  to  two  complete  sets  of 
compounds.  The  nomenclature  used  to  distinguish  the  two  series 
has  been  discussed  before  (p.  70).  As  a  rule,  an  element  passes  from 
one  form  of  combination  to  another  without  change  of  valence.  But 
compounds  of  elements  like  tin  can  also  undergo  changes  in  course 


106  INORGANIC   CHEMISTRY 

of  which  the  valence  alters.     Cases  of  this  kind  will  be  considered 
•when  they  arise  (see  Preparation  of  chlorine). 

The  regular  valence  of  an  element  cannot  be  learned  by  examining  the  com- 
position of  a  compound  chosen  at  random.  Thus  FeS,  H2S,  HgS,  and  other  com- 
pounds show  sulphur  to  be  bivalent.  There  is  also  a  series  in  which  sulphur  is 
sexivalent,  as  in  SO3.  But  the  compound  S2O3,  in  which  sulphur  appears  to  be 
trivalent,  is  an  isolated  case.  Again,  FeO,  FeS,  FeCl2  show  iron  to  be  bivalent, 
and  FeCl3,  Fe2(SO4)3,  etc.,  show  it  to  be  also  trivalent.  But  Fe3O4,  the  magnetic 
oxide,  is  an  exception.  Valence  has  to  do  mainly  with  chemical  interactions,  in 
which  the  element  either  passes  from  one  state  of  combination  to  another  without 
change  of  valence,  or  goes  over  into  a  compound  of  another  regular  series  with 
another  regular  valence.  It  is  not  a  matter  of  statics.  Hence,  questions  as  to  the 
magnitude  of  the  valence  in  isolated  compounds  like  Fe304,  N2O,  and  so  forth,  are 
at  present  of  minor  importance. 

A  definition  of  valence  differing  from  those  given  above  is  preferred  by  many 
chemists.  The  atomic  weight  of  a  univalent  element  can  hold  but  one  unit  of  an- 
other element  in  combination.  Thus,  the  weight  of  chlorine  represented  by  Cl  can 
hold  but  one  H  or  one  Na  in  combination.  An  atomic  weight  of  a  bivalent  ele- 
ment, although  it  combines  with  but  one  unit  of  another  bivalent  element,  may 
hold  as  many  as  two  units  of  a  univalent  element  in  combination.  But  it  cannot 
hold  more.  A  unit  of  a  trivalent  element,  however,  may  hold  as  many  as  three 
units,  provided  the  other  element  is  univalent.  In  this  point  of  view  the  valence 
of  an  element  is  the  maximum  capacity  of  its  atomic  weight  to  hold 
atomic  weights  of  other  elements  in  combination. 

Valence  is  often  defined  as  the  power  of  the  atomic  weight  of  one  element,  to 
hold  units  of  other  elements  in  combination.  But  the  word  power  suggests  that 
valence  is  a  measure  of  the  force  with  which  the  elements  are  held  together, 
whereas  it  has  to  do  with  the  quantity  of  matter  only.  Gold  is  trivalent,  but  holds 
chlorine  with  incomparably  less  force  than  does  sodium  which  is  only  univalent. 

Physical  Properties  of  Hydrogen.  —  Some  of  these  may  be 
given  in  tabular  form  : 

Colorless  Crit.  temp.,  about  —  234° 

Tasteless  Sp.  Ht.  (gas),  3.4 

Odorless  Boiling-point,  —  252.5° 

Density  (air  =  1),  0.0695          Melting-point  (58  mm.),-  260° 
Density  (H  =  1),  1  Sol'ty  in  Aq,  1.9  vols.  in  100  (14°) 

Wt.  of  1  1.,  0.08987  g. 

Air  is  14.5  times  as  heavy,  hence  the  gas  may  be  poured  upwards 
and  is  used  for  filling  balloons.  A  liter  flask  filled  with  air  requires 
about  1.2  g.  to  be  added  to  the  tare  to  restore  the  balance  when  the 
air  is  displaced  by  hydrogen.  Its  specific  heat  is  about  seventeen 
times  that  of  oxygen  (0.2).  Its  thermal  conductivity  is  greater  than 
that  of  any  other  gas.  Hence  a  wire,  raised  to  incandescence  in 


HYDROGEN  107 

air  by  means  of  an  electric  current,  cannot  be  kept  at  a  red  heat,  even, 
by  the  same  current  in  hydrogen. 

Hydrogen  was  first  liquefied  in  visible  amounts  by  Dewar  (1898). 
The  liquid  is  colorless,  and,  when  allowed  to  evaporate  rapidly  under 
reduced  pressure,  freezes  to  a  colorless  solid.  All  other  gases  except 
helium  solidify  easily  when  led  into  a  vessel  surrounded  by  liquid 
hydrogen. 

Hydrogen  is  absorbed,  for  the  most  part  in  a  purely  mechani- 
cal way,  by  many  metals.  Heated  iron  will  take  up  19  times  its 
volume  of  hydrogen.  Under  similar  conditions  gold  takes  up  46  vol- 
umes, platinum  in  fine  powder  50  volumes,  palladium  502  volumes, 
and  silver  none.  The  maximum  absorbed  by  palladium  under  favor- 
able conditions  is  873  volumes.  It  is  still  a  question  whether,  in  the 
case  of  palladium,  a  part  of  the  gas  is  not  in  combination. 

Diffusion.  —  If  a  volume  of  gas  is  inclosed  at  one  end  of  a  cylin- 
der, the  rest  of  which  is  entirely  empty,  and  is  suddenly  released  from 
this  confinement,  it  spreads  with  extreme  speed  so  as  to  occupy  the 
whole  of  the  cylinder  to  an  equal  degree.  This  spreading  is  not  an 
effect  of  gravitation,  since  it  takes  place  upwards  or  downwards 
with  equal  celerity.  The  same  phenomenon  is  observed  when,  in  every- 
day life,  a  bottle  of  scent  is  opened.  The  vapor,  on  escaping,  begins 
to  penetrate  in  all  directions  through  the  room,  showing  its  presence 
by  its  odor.  The  motion,  as  this  instance  shows,  takes  place  through 
a  space  occupied  by  another  gas  more  slowly  than,  but  just  as  surely 
as,  when  the  space  is  empty.  The  material  of  gases  has  in  fact  an  in- 
dependent power  of  locomotion.  The  resulting  phenomenon  we  call 
diffusion.  It  is  constant  in  rate  for  each  gas  under  like  conditions, 
and  hydrogen  has  the  greatest  speed  of  diffusion  of  all  the  gases. 

The  interdiffusion  of  gases  and  the  absence  of  gravity  effect  may  be  shown 
simultaneously.  A  jar  is  filled  with  carbon  dioxide  and  a  jar  of  air  is  inverted 
and  placed  mouth  to  mouth  with  the  other.  After  a  few  minutes,  and  in  spite  of 
the  fact  that  carbon  dioxide,  measured  in  bulk,  is  one-half  heavier  than  air,  as 
much  of  each  gas  will  be  found  in  the  cylinder  of  the  other  as  in  its  own.  Lime- 
water  (p.  68)  will  show  the  presence  of  carbon  dioxide  in  the  upper  jar.  The  phe- 
nomena of  diffusion  must  not  be  confused  with  cases  like  the  pouring  of  hydrogen 
upward  to  displace  air  in  an  inverted  jar.  In  this  case  the  gas  flows  en  masse,  and 
the  gravity  effect  is  the  very  one  on  which  we  depend  for  the  success  of  the  experi- 
ment. It  is  when  hydrogen  scatters  itself  in  a  somewhat  slower  way,  and  down- 
ward and  sideways  as  well  as  upward,  that  we  have  diffusion.  The  word  indicates 
the  scattering  rather  than  the  flowing  nature  of  the  phenomenon. 


108 


INORGANIC   CHEMISTRY 


The  different  rates  of  diffusion  of  different  gases  are  easily  shown 
by  comparing  their  several  speeds  with  that  of  air,  when  both  pass 
through  a  wall  of  unglazed,  porous  porcelain. 

The  porous  cylinder  A  on  the  left  (Fig.  38)  contains  air  and  is  con- 
nected with  a  wide  tube  which  dips  beneath  the  surface  of  the  water. 
When  a  cylinder  If  containing  hydrogen  is  brought  over  it,  rapid 
escape  of  gas  takes  place  through  the  water,  showing  that  a  rise  in 
pressure  has  taken  place  inside  the  porous  vessel.  Before  the  cylinder 
of  hydrogen  approached  it,  the  air  was  moving  both  outwards  and  in- 
wards through  the  porcelain,  but, 
being  the  same  air,  the  speed  of 
motion  was  equal  in  both  direc- 
tions, and  therefore  the  pressure 
inside  was  not  affected.  It  is  im- 
portant to  note  that  there  was  at 
no  time  rest,  there  was  simply 
equal  motion  in  both  directions. 
When  the  hydrogen  atmosphere 
surrounded  the  cylinder,  the  hy- 
drogen gas  moved  more  rapidly 
into  the  cylinder  than  the  air  in- 
side could  move  out,  and  hence  an 
excess  of  pressure  quickly  arose 
in  the  interior. 

The  cylinder  on  the  right  is 
similar,  but  a  vessel  C  filled  with 
carbon  dioxide,  a  gas  heavier  than 
air  (density  1.53,  air  =  1),  sur- 
rounds it.  Here  the  air  moves 
out  faster  than  the  gas  can  move 
in,  a  reduction  in  pressure  takes 
place,  and  the  water  rises. 
Exact  measurement  shows  that  the  lighter  a  gas  is  in  bulk,  the 
faster  its  parts  move  by  diffusion  in  any  direction.  The  rate  is 
inversely  proportional  to  the  square  root  of  the  density  of  the  gas. 
Thus,  for  hydrogen  and  air  it  is  in  the  ratio  Vl  :  V.0695,  or  3.8  :  1. 
For  air  and  carbon  dioxide  it  is  Vl.53  :  VI,  or  1.24  :  1. 

Chemical  Properties  of  Hydrogen.  —  Hydrogen,  delivered  from 
a  jet,  burns  in  air  or  pure  oxygen.  A  cold  vessel  held  over  the  almost 


FIG.  38. 


HYDROGEN  109 

invisible  blue  flame  condenses  to  droplets  of  water  the  steam  that  is 
produced.  Although  the  flame  gives  little  light,  it  is  exceedingly  hot. 
Platinum  melts  in  it  easily.  In  a  closed  space  it  produces  a  tem- 
perature of  over  2500°.  When  hydrogen  and  oxygen  are  mingled  in  a 
suitable  burner,  and  the  flame  is  allowed  to  play  on  a  piece  of  quick- 
lime, the  latter  becomes  white-hot  at  the  spot  where  the  flame  meets 
it.  This  result  is  called  a  calcium  light  or  lime  light. 

When  the  gases  are  mixed  in  a  glass  vessel,  the  chemical  action  is 
very  slow  at  ordinary  temperatures,  no  perceptible  amount  of  union 
occurring  in  a  period  of  five  years.  If  the  mixture  is  sealed  up  and 
kept  at  300°,  after  several  days  a  small  part  is  found  to  have  com- 
bined to  form  water.  At  518°,  hours  are  required  before  the  union  is 
complete.  At  600°  the  interaction  is  rapid,  but  not  explosive.  At 
700°  the  combination  is  almost  instantaneous.  Hence  contact  with 
a  body  at  a  bright-red  heat  (p.  73)  is  required  actually  to  explode  the 
mixture. 

These  facts  illustrate  the  effect  of  temperature  on  the  speed  of 
chemical  changes  (p.  72).  A  rough  calculation  shows  that,  since 
interactions  lower  their  speed  to  half  its  value  for  every  depression  of 
10°  in  temperature,  at  ordinary  temperatures  this  union  can  hardly 
make  easily  perceptible  progress  in  less  than  a  thousand  million  years. 
This  effect  of  temperature,  therefore,  accounts  for  the  apparent  absence 
of  action  in  the  cold  gases. 

Finely  divided  platinum,*  when  held  in  the  mixture,  hastens  the 
action  in  the  part  of  the  gases  in  contact  with  it.  'The  heat  of  their  union 
raises  the  temperature  of  neighboring  portions  and  causes  explosion  of 
the  mass.  The  platinum  is  simply  a  catalytic  agent  (p.  75)  and 
remains  itself  unaffected.  Its  role  is  to  increase  prodigiously  the 
vanishingly  small  speed  of  the  union  between  the  cold  gases. 

Hydrogen  unites  directly  with  a  minority  only  of  the  simple  sub- 
stances. It  combines  rapidly  with  oxygen,  chlorine,  fluorine,  and 
lithium,  and  more  slowly  with  a  few  others. 

When  these  elements,  especially  the  first  two,  are  already  in  com- 
bination, hydrogen  may  still  sometimes  displace  the  material  with 
which  they  are  united.  Thus,  when  one  of  the  oxides  of  copper  or  of 
iron  is  heated  in  a  tube  through  which  hydrogen  flows,  the  latter  com- 

*  The  most  convenient  form  is  obtained  by  dipping  asbestos  in  a  solution  of 
chloroplatinic  acid,  and  heating  it  in  the  blast-lamp.  The  fibers  are  covered  with 
a  thin  film  of  the  metal : 

H,PtCl4-»  Pt  +  4C1  +  2HC1 . 


110  INORGANIC   CHEMISTRY 

bines  with  the  oxygen  to  form  water,  and  the  metal  is  liberated.  To 
make  the  equations,  we  set  down  first  the  formulae  of  the  substances 
used  and  produced  in  each  case  : 


H->H20  +  Cu, 
Fe3O4-f-H->H20  +  Fe. 

We  then  observe  that,  for  each  atomic  weight  of  oxygen,  2H  will  be 
required,  and  amend  the  equations  thus  : 

CuO-f  2H-*H20  +  Cu,  (1) 

Fe8O4  +  8H  ->  4H20  +  Fe. 

Then  we  make  the  amount  of  iron  produced  equal  to  that  taken  : 

Fe804  +  8H  -»  4H2O  +  3Fe.  (2) 

Thus  (1)  and  (2)  are  the  final  equations. 

These  interactions  are  classed  as  displacements.  In  describing 
them  the  chemist  would  also  say  that  the  hydrogen  has  been  oxidized 
and  that  the  oxide  of  the  metal  has  been  reduced  (p.  72). 

An  Inapt  Use  of  the  Word  "Affinity  "  in  Explanation  of 
Chemical  Actions.  —  It  has  passed  into  the  common  language  of 
chemistry  that  actions  like  the  reduction  of  magnetic  oxide  of  iron, 
just  mentioned,  are  "  explained  "  by  saying  that  the  hydrogen  has  a 
greater  tendency  to  unite  with  oxygen,  or  has  a  greater  affinity  for  it, 
than  has  iron,  and  therefore  removes  the  oxygen  from  combination 
with  the  latter.  Plausible  as  this  statement  seems,  it  would  be  in 
most  cases,  as  here,  quite  incorrect.  Under  the  modes  of  preparing 
hydrogen,  we  spoke  of  the  action  of  steam  upon.  iron  (p.  99),  and  gave 
the  equation  :  3Fe  +  4H20  —  >  Fe804  -f-  8H.  To  use  consistently  this 
handy  method  of  explaining  chemical  change  by  the  help  of  the  word 
"  affinity,"  we  should  have  to  say  that  the  hydrogen  has  a  less  amnity 
for  the  oxygen  than  has  iron,  and  therefore  hydrogen  is  set  free  and 
oxide  of  iron  is  formed.  It  will  be  seen  that  this  statement  is  in  direct 
contradiction  to  the  one  made  above.  Both  cannot  be  true.  The  fact 
is  that  both  are  based  upon  an  assumption  which  is  incorrect  —  the 
assumption,  namely,  that  the  displacement  of  one  element  by  another 
is  always  an  evidence  of  the  greater  affinity  of  the  latter.  Until  the 
means  of  truly  measuring  affinity,  or,  as  we  prefer  to  call  it,  activity, 
have  been  explained,  we  shall  do  well,  as  far  as  possible,  to  avoid  using 
the  word. 


HYDROGEN  111 

The  action  of  catalytic  agents  is  itself  a  refutation  of  this  blundering  assump- 
tion. Putting  a  little  platinum  in  a  mixture  of  oxygen  and  hydrogen  cannot  add 
to  the  energy  contained  in  these  substances,  and  cannot  therefore  increase  their 
intrinsic  tendencies  to  unite.  Yet  in  its  presence  an  almost  nonexistent  action 
becomes  suddenly  explosively  violent.  There  are  other  far  more  potent  factors 
than  affinity  which  determine  the  direction  and  speed  of  many  chemical  changes 
(see  Chemical  equilibrium). 

In  this  connection,  it  is  worth  noting  that,  while  increasing  the  speed  of  a  train 
or  a  ship  requires  a  great  addition  to  the  energy  expended,  and  is  very  costly,  in- 
creasing the  speed  of  a  chemical  change  requires  the  expenditure  of  no  energy 
whatever.  The  employment  of  a  couple  (p.  96)  or  a  catalytic  agent  adds  nothing 
to  the  energy  the  separate  bodies  possessed  before  they  were  mixed.  And  the  cata- 
lytic agent  is  recovered  unchanged  and  as  efficient  as  ever  at  the  end.  Theoretically, 
therefore,  these  agencies  cost  nothing.  The  increased  speed  in  the  formation  of 
the  products  is  obtained  gratis.  The  contact  method  of  making  sulphuric  acid 
(g.u.)  illustrates  the  way  in  which  commerce  has  taken  advantage  of  this  fact. 

The  Speed  of  Chemical  Actions :  a  Means  of  Measuring 
Activity.  —  The  speed  of  a  chemical  action  is  measured  by  the  num- 
ber of  atomic  or  formula  weights  of  the  substance  undergoing  change 
in  a  given  time.  Now,  one  means  of  measuring  the  relative  chemical 
activities  of  several  substances,  and,  therefore,  of  the  relative  amounts 
of  available  chemical  energy  they  contain  (p.  26),  is  to  observe  the 
speed  with  which  they  undergo  the  same  chemical  change  (p.  28). 
Thus  we  may  compare  the  activities  of  the  various  metals  by  allowing 
them  separately  to  interact  with  hydrochloric  acid  and  collecting  and 
measuring  the  hydrogen  liberated  per  minute  by  each.  It  will  be  seen, 
even  in  the  roughest  experiment,  that  magnesium  is  thus  much  more 
active  than  zinc.  The  comparison  must  be  made  with  such  precau- 
tions, however,  as  will  make  it  certain  that  the  conditions  under  which 
the  several  metals  act  are  all  alike.  Thus,  in  spite  of  the  heat  evolved 
by  the  action,  means  must  be  used,  by  suitable  cooling,  to  keep  the 
temperature  at  some  fixed  point  during  the  experiment,  for  all  actions 
become  more  rapid  when  the  temperature  rises  (p.  72).  Again,  the 
pieces  of  the  various  metals  must  be  arranged  so  that  equal  surfaces 
are  exposed  to  the  acid  in  each  case ;  for  pieces  of  the  same  metal, 
having,  of  course,  the  same  intrinsic  activity,  will  nevertheless  give 
hydrogen  more  rapidly  the  larger  the  surface  they  expose.  Equal 
weights  of  zinc  will  finally  give  equal  weights  of  hydrogen ;  but  if  one 
of  them  is  in  the  form  of  foil  while  the  other  is  a  cylinder,  the  former, 
although  it  will  not  last  so  long,  will  give  much  more  hydrogen  per 
minute.  Still  again,  the  portions  of  hydrochloric  acid  must  contain 
the  same  percentage  of  hydrogen  chloride  in  each  case,  for  the  metal 


112  INORGANIC   CHEMISTRY 

will  secure  the  acid  it  needs  with  less  delay  in  a  more  concentrated 
solution  than  in  a  less  concentrated  solution,  and  in  the  former  case 
will  therefore  displace  hydrogen  more  rapidly.  When  these  and  other 
precautions  have  been  taken,  a  true  comparison  of  the  relative  activi- 
ties of  the  metals  with  respect  to  this  particular  action  may  be  made. 
It  is  found  that  the  order  in  which  this  comparison  places  the  metals 
is  much  the  same  as  that  in  which  they  are  placed  by  a  study  of  other 
similar  actions.  This  is  natural,  since  we  are  really  comparing,  in  each 
case,  the  amount  of  chemical  energy  in  each  metal.  A  single  table 
suffices,  therefore,  for  all  purposes.  This  table  is  given  in  connection 
with  a  more  exact  method  of  comparing  the  activities  of  the  metals 
(see  Electromotive  series  of  the  metals).  The  speed  of  chemical 
changes  is  discussed  in  greater  detail  under  Chemical  Equilibrium  (see 
also  Sulphurous  acid). 

Exercises.  —  1.  What  are  the  valences  of  the  negative  radicals 
of  phosphoric  acid  (p.  71),  and  of  acetic  acid  (p.  93)  ?  What  must  be 
the  formulae  of  calcium  phosphate,  cupric  acetate,  aluminium  phos- 
phate, ferrous  carbonate,  ferrous  sulphate,  cupric  chloride  ? 

2.  What  is  the  valence  of   phosphorus   in  phosphoric   anhydride 
(p.  71)  ?     What  must  be  the  formulae  of  the  chloride  and   the   sul- 
phide of  phosphorus,  and  of  aluminium  oxide  ? 

3.  What  are  the  valences  of  the  elements  in. the  following:  LiH, 
NH8,  SeH2,  BN? 

4.  WThat  are  the  valences  of  the  metals  and  radicals  in  the  following : 
Pb(N08)2,  Ce(S04)2,  KC1,  KMn04  (potassium  permanganate)  ?     Name 
all  the  substances  in  3  and  4. 

5.  Write  the  formulae  of  ferrous  and  ferric  oxides,  of  ferrous  and 
ferric  nitrates,  of  stannous  and  stannic  sulphides. 

6.  What  must  be  the  relative  rates  of  diffusion  of  hydrogen  and 
of  carbon  dioxide  ? 

7.  Make  equations  to  represent  (a)  the  reduction  of  lead  dioxide 
(Pb02)  by  hydrogen ;  (b)  the  actions  of  aluminium  upon  cold  water 
and  upon  steam  at  a  red  heat. 


CHAPTER  VIII 
WATER 

THE  great  quantity  of  water  which,  occurs  in  nature  makes  it  one  of 
the  most  familiar  chemical  substances.  The  ocean  covers  about  three- 
fourths  of  the  surface  of  the  earth,  and  in  most  habitable  regions  lakes 
and  streams  abound.  Water  is  found  also  in  the  bodies  of  both  ani- 
mals and  plants  in  large  quantities,  and  is  indeed  essential  to  the 
working  of  living  organisms. 

Natural  Waters.  —  The  water  found  in  nature  varies  greatly  in 
the  amount  of  foreign  material  which  it  contains.  Sea-water  holds 
about  3.6  per  cent  of  solid  matter  in  solution,  while  rain-water  is  the 
purest  natural  water.  Even  rain-water  contains  foreign  matter,  how- 
ever. When  we  heat  it,  bubbles  of  gas  form  on  the  sides  of  the  vessel, 
showing  that  oxygen  and  nitrogen  from  the  air  have  been  dissolved  by 
the  water  as  it  fell.  On  evaporating  a  considerable  mass  of  such 
water,  we  find  that,  aside  from  dust,  crystals  of  chemical  substances, 
such  as  ammonium  nitrate,  may  be  recognized  in  the  residue.  Of  well 
and  surface  waters,  some  which  contain  calcium  sulphate,  calcium 
bicarbonate,  and  compounds  of  magnesium  in  solution  are  described  as 
hard.  Others  contain  compounds  of  iron,  and  still  others  are  effer- 
vescent and  give  off  carbon  dioxide.  These  are  called  mineral  waters. 
All  of  the  dissolved  substances  are  obtained  by  the  water  in  its  prog- 
ress over  or  under  the  surface  of  the  ground. 

Water  which  is  to  be  used  for  domestic  purposes  is  examined,  not 
only  to  ascertain  the  amount  of  the  ingredients  which  produce  hard- 
ness, but  also  with  reference  to  the  proportion  of  organic  matter  which 
it  may  hold  in  solution.  This  usually  gains  access  to  the  water  by 
admixture  of  sewage  (p.  72).  It  is  not  the  organic  matter  itself  which 
is  deleterious,  but  the  bacteria  of  putrefaction  and  disease  which  are 
likely  to  accompany  it.  Inoculation  of  culture  media  with  the  water 
can  alone  show  whether  or  not  the  latter  are  present. 

Purification  of  Water.  — The  foreign  materials  which  water  may 
contain  are  divisible  into  two  kinds, — dissolved  matter  and  suspended 

113 


114  INORGANIC   CHEMISTRY 

matter.  No  water  is  free  from  either  of  these  varieties  of  impurity. 
In  chemical  laboratories  distilled  (p.  38)  water  of  a  more  or  less  pure 
kind  is  always  employed,  but  this  represents  only  a  crude  purification. 
By  using  a  platinum  still  and  condenser,  water  of  much  greater  purity 
can  be  obtained.  Yet,  on  account  of  the  solvent  power  of  water,  it  is 
impossible  to  keep  such  a  liquid  even  for  a  short  time.  Ordinary 
glass  dissolves  in  water  to  a  very  noticeable  extent. 

The  purity  of  water  is  most  easily  investigated  by  measurement  of 
its  resistance  to  the  passage  of  electricity.  A  column  only  one  milli- 
meter long,  of  the  purest  distilled  water  that  can  be  made,  has  a  greater 
resistance  than  a  copper  wire  of  the  same  cross-section  and  long  enough 
to  reach  a  thousand  times  round  the  earth  at  the  equator.  During  a 
few  minutes'  exposure  to  the  air,  or  contact  with  a  glass  vessel,  how. 
ever,  a  sufficient  amount  of  foreign  material  (p.  95)  of  high  conductiv- 
ity is  taken  up  to  diminish  its  resistance  very  greatly. 

For  ordinary  purposes  the  suspended  matter  which  water  contains 
is  removed  by  filtration  (p.  11).  In  the  laboratory  this  takes  place 
through  unsized  paper.  The  pores  of  the  paper  are  sufficiently  small 
to  retain  particles  of  the  dimensions  usually  met  with,  while  permit- 
ting the  passage  of  the  water  with  its  dissolved  matter.  On  a  large 
scale,  beds  of  gravel  are  employed.  In  the  household  the  Pasteur 
filter  is  more  compact  and  efficient.  The  water  is  forced  by  its  OWE 
pressure  through  the  pores  of  a  closed  tube  made  of  unglazed  porce- 
lain. Care  must  be  taken  to  clean  these  tubes  at  frequent  intervals, 
so  that  organic  and  perhaps  putrescent  matters  may  not  accumulate 
upon  them.  If  the  cleansing  is  not  carried  out,  the  Pasteur  filter-tube 
becomes  a  breeding-place  for  bacteria,  and  may  add  greatly  to  the  con- 
tamination of  the  water  instead  of  diminishing  it. 

Matter  in  solution  cannot  be  removed  by  filtration,  and  is  eliminated 
by  distillation.  Since  the  water  is  converted  into  steam  and  is  con- 
densed in  platinum  or  tin  pipes,  only  gases  or  volatile  liquids  dis- 
solved in  it  can  pass  into  the  distillate. 

Physical  Properties  of  Water.  — When  we  view  a  white  object 
through  a  deep  layer  of  water  we  find  that  the  liquid  has  a  blue  or 
greenish-blue  color.  At  a  pressure  of  760  mm.,  it  exists  as  a  liquid 
between  0°  and  100°.  Below  0°  it  becomes  solid,  above  100°  a  gas.  Of 
all  chemical  substances  it  is  the  .one  which  we  use  most,  so  that  famil- 
iarity with  its  properties  is  indispensable  to  the  chemist.  It  will  serve 
also  as  a  typical  liquid,  since  it  differs  from  others  only  in  details. 


WATER  115 

The  weight  of  a  cubic  centimeter  of  water  at  4°  gives  us  our  unit,  the  gram.  A 
kilogram  of  water  at  0°  occupies  1.00013  liters,  or  0.13  c.c.  more  than  at  4°  C.  A 
kilogram  of  ice  at  0°  occupies  1.09083  liters,  or  90.7  c.c.  more  than  an  equal  weight 
of  water.  The  volume  of  the  same  weight  of  water  at  100°  is  1.0432  liters. 

Ice.  —  The  raising  or  lowering  of  the  temperature  of  a  gram  of 
water  through  one  degree  corresponds  to  the  addition  or  removal  of 
one  calorie  of  heat.  The  conversion,  however,  of  a  gram  of  water  at 
0°  to  a  gram  of  ice  at  0°  requires  the  removal  of  79  calories  of  heat. 
The  mere  melting  of  a  gram  of  ice  causes  an  absorption  of  heat  to  the 
same  amount.  This  is  called  the  heat  of  fusion  of  ice.  At  0°a  mix- 
ture of  ice  and  water  will  remain  in  unchanged  proportions  indefinitely. 
Any  cause  which  tends  permanently  to  lower  or  raise  the  temperature 
by  a  fraction  of  a  degree,  however,  will  bring  about  the  disappearance 
of  the  water  or  of  the  ice  respectively.  This  temperature  is  called 
the  melting  or  the  freezing  point.  Properties  of  this  kind,  marked 
by  transition  points  from  one  state  to  another,  are  much  used  in  chem- 
istry for  keeping  other  bodies  or  systems  at  a  constant  temperature 
during  measurement  or  observation.  A  mixture  of  ice  and  water  sur- 
rounding a  body,  when  kept  in  constant  agitation,  will  automatically 
maintain  the  body  at  a  fixed  temperature  (0°)  so  long  as  both  compo- 
nents hold  out. 

Steam,  and  Aqueous  Tension.  —  At  atmospheric  pressure,  water 
passes  into  steam  rapidly  at  100°,  but  at  lower  temperatures,  and  even 
when  frozen,  it  does  the  same  thing  more  slowly.  The  best  way  to 
define  the  quantity  of  the  vapor  at  various  temperatures  is  by  the 
gaseous  pressure  it  exercises.  This  is  proportional  to  the  concentra- 
tion (p.  81)  of  the  aqueous  material  in  the  space  above  the  water, 
and  has  a  definite  value  for  each  temperature.  Its  amount  may  be 
shown  by  allowing  a  few  drops  of  water  to  ascend  into  the  vacuum  at 
the  top  of  a  barometric  column  (Fig.  39).  The  tube  on  the  left  shows 
the  mercury  when  nothing  presses  on  its  surface.  The  tube  on  the 
right  shows  the  result  of  admitting  the  water.  The  tension  of  the 
atmosphere  being  the  same  for  both,  the  smaller  height  of  mercury 
which  now  suffices  to  counterbalance  it  shows  that  something,  which 
can  be  nothing  but  the  water  vapor,  is  pressing  on  the  surface  of  the 
mercury,  and  makes  up  the  rest  of  the  total  stress  needed.  The  differ- 
ence in  the  height  of  the  two  columns  gives  the  value  of  this  pressure, 
which  we  call  the  vapor  pressure  of  the  water.  The  jacket  surround- 
ing the  tube  on  the  right  enables  us,  by  adding  ice  or  warm  water,  to 


116 


INORGANIC   CHEMISTRY 


keep  the  water  that  is  admitted  to  the  vacuum,  and  the  parts  of  the 
apparatus  immediately  in  contact  with  it,  at  any  temperature  between 
0°  and  100°. 

When  ice  is  used  outside,  and  a  piece  of  it  is  introduced  into  the 
vacuum,  the  vapor  it  gives  off  quickly  reaches  a  pressure  of  4.5  mm. 
The  vapor  pressure  of  the  ice  takes  the  place  of  4.5  mm.  of  mercury  in 

balancing  the  atmospheric  pressure,  and 
so  the  mercury  column  falls  by  this 
amount.  Similarly,  water  at  10°  causes 
a  fall  of  9.1  mm.  and  at  20°  of  17.4  mm., 
so  that  these  represent  the  mercury- 
height  values  of  the  vapor  pressure  at 
these  temperatures.  The  quantity  of 
water  used  makes  no  difference,  so  long 
as  a  little  more  is  present  than  is  re- 
quired to  fill  the  available  space  with 
vapor.  Of  course,  if  a  large  amount 
is  admitted,  its  dead  weight  will  take 
the  place  of  an  equal  weight  of  mer- 
cury in  balancing  the  pressure  of  the 
air.  If  there  is  a  measurable  column 
of  water,  its  height  must  be  divided 
by  13.6  (the  sp.  gr.  of  mercury),  and 
counted  as  if  it  were  part  of  the  mercury. 
An  expression  is  needed  to  describe 
the  varying  power  of  the  water  at 
different  temperatures  to  maintain  dif- 
ferent pressures  of  its  vapor.  This  we 
call  the  aqueous  tension  of  the  liquid. 
Its  magnitude  at  any  given  temperature 
is  measured  by  the  maximum  pressure 

reached  by  the  vapor  (see  Kinetic-molecular  hypothesis  applied  to 
liquids). 

With  water  at  higher  temperatures  the  fall  of  the  mercury  column 
becomes  much  greater.  At  50°  it  is  92  mm.,  at  70°  it  is  233.3  mm.,  at 
90°  it  is  525.5  mm.,  and  at  100°  it  is  760  mm.  At  the  boiling-point, 
therefore,  the  aqueous  tension  takes  the  place  of  the  whole  barometric 
column,  and  is  equal  to  the  average  air  pressure.  At  121°  the  aqueous 
tension  is  two  atmospheres,  at  180°  it  is  ten  atmospheres. 

There  is  another  standpoint  from  which  these  phenomena  may  be 


FIG.  39. 


WATER  117 

viewed.  Water  vapor  can  exist  at  10°  only  if  the  pressure  upon  it  is 
9.1  mm.  or  less.  If  we  imagine  the  water  placed  in  a  cylinder  closed 
by  a  frictionless,  weightless  piston  (Fig.  40),  then  at  10°  the  piston 
will  remain  at  rest  whether  we  place  it  high  or  low,  provided  it  is 
loaded  with  a  weight  exactly  equal  to  that  of  a  layer  of  mercury  9.1  mm. 
thick  covering  its  whole  area.  We  speak  of  such  a  system  as  being  in 
equilibrium.*  With  a  less  weight  the  piston  will  move  slowly  upwards, 
as  the  vapor  continually  given  off  by  the  water  presses  upon  it,  until 
it  reaches  the  top  or  the  water  all  evaporates.  Conversely, 
if  it  bears  a  greater  load,  it  will  move  down  and  the  vapor 
will  condense  on  the  walls  and  bottom  of  the  cylinder  until 
the  piston  comes  in  contact  with  the  water  itself  and  the 
vapor  is  all  abolished.  These  conceptions  will  find  con- 
stant application  not  only  to  physical  but  also  to  chemical 

phenomena  (see  Kinetic-molecular  hypothesis).     The  ex-    I L 

pression : 

Aq.  (liq.)  <=±  Aq.  (vap.) 

is  used  to  represent  the  state  of  equilibrium  in  a  system  like      PIG.  40. 
the  above. 

One  other  phase  of  this  subject  is  recognized  by  special  phraseology. 
When  water  at  a  certain  temperature  has  given  the  full  amount  of 
water  vapor  to  the  space  above  it  that  its  aqueous  tension  permits,  we 
say  that  the  space  is  saturated  with  vapor.  That  concentration  of 
vapor  which  constitutes  saturation  varies  with  the  temperature  of  the 
water  and  depends  therefore  solely  on  the  power  of  the  water  to  give 
off  vapor.  It  has  nothing  to  do  with  the  size  of  the  space,  and  is  even 
independent  of  other  gases  the  space  may  already  contain  (p.  88). 

The  space  immediately  above  the  surface  of  the  ground,  which  is 
mainly  occupied  by  atmospheric  air,  is,  on  an  average,  less  than  two- 
thirds  saturated  with  water  vapor.  That  is  to  say,  such  air,  when 
inclosed  in  a  vessel  containing  water,  will  take  up  about  one-half  more 
than  it  .already  contains.  At  100°  the  water  vapor  displaces  the  air 
entirely,  and  the  liquid  is  said  to  boil. 

The  water  present  in  the  air  plays  an  important  part  in  many 
chemical  phenomena,  as  we  shall  see.  All  our  substances  and  appara- 
tus have  traces  of  water  condensed  on  their  surfaces.  This  water  is, 
in  a  sense,  in  an  abnormal  condition,  for  it  does  not  evaporate  even  in 

*  In  chemistry  we  do  not  speak  of  stable,  unstable,  and  indifferent  equilib- 
rium, as  is  done  in  physics,  but  only  of  the  first.  The  term  "  inetastable,"  how- 
ever, is  employed  (see  Chap.  x). 


118  INORGANIC   CHEMISTRY 

dry  air.  It  is  observed  to  pass  off  in  vapor,  however,  when  we  have 
occasion  to  heat  any  material. 

In  passing  into  vapor,  water  absorbs  heat  without  changing  its 
temperature.  A  gram  of  water  at  100°,  for  example,  in  turning  into  a 
gram  of  steam  at  100°,  takes  up  537  calories.  This  is  called  its  heat 
of  vaporization.  Steam,  in  fact,  contains  much  more  internal  energy 
than  an  equal  weight  of  water  at  the  same  temperature,  just  as  water, 
in  turn,  contains  more  energy  than  ice. 

The  temperature  of  100°  is,  like  the  melting-point  of  ice,  an  im- 
portant transition  point,  and  has  the  same  properties,  mutatis  mutan- 
dis, as  the  latter.  It  is  less  exactly  recoverable  by  simply  keeping  a 
vessel  full  of  water  in  ebullition,  however,  because  variations  in  the 
pressure  of  the  atmosphere  affect  it  more  markedly  than  they  do  the 
melting-point  of  ice.  Near  to  100°,  the  boiling-point  rises  or  falls 
about  0.037°  for  1  mm.  change  in  pressure  (cf.  p.  116).  On  the  top 
of  Mont  Blanc  water  boils  at  84°. 

Water  as  a  Solvent.  —  One  of  those  physical  properties  of  water 
which  are  most  used  in  chemical  work  is  its  tendency  to  dissolve  many 
substances.  This  subject  is  so  important  and  extensive  that  we  shall 
presently  devote  a  complete  chapter  to  some  of  its  simpler  and  more 
familiar  aspects. 

Chemical  Properties  of  Water.  —  Water  is  so  very  frequently 
used  in  chemical  experiments  in  which  it  is  a  mere  mechanical  ad- 
junct, that  the  beginner  has  difficulty  in  distinguishing  the  cases  in 
which  it  has  itself  taken  part  in  the  chemical  interaction.  The  rather 
limited  list  of  kinds  of  chemical  activity  it  can  show  should  therefore 
receive  careful  notice :  It  (1)  is  a  relatively  stable  substance.  It  (2) 
combines  directly  with  many  substances.  In  the  commoner  of  the  two 
kinds  of  this  action,  the  formation  of  hydrates  (see  below),  the  com- 
pound exists  in  the  solid  form  only,  however,  and  is  decomposed  in 
solution.  Finally,  it  (3)  interacts  with  some  substances  in  a  way  which 
we  describe  as  hydrolysis.  We  shall  not  discuss  this  last  property 
until  some  substance  which  is  markedly  affected  is  encountered  (see 
Preparation  of  hydrogen  chloride). 

Perhaps  we  should  add  that  steam  at  a  high  temperature  oxidizes 
elements  which  readily  combine  with  oxygen.  For  example,  it  turns 
iron  into  the  magnetic  oxide  (p.  99).  At  such  high  temperatures, 
however,  the  water  is  partially  resolved  into  a  mixture  of  hydrogen  and 


WATER  119 

oxygen,  and,  the  latter  being  the  more  active  of  the  two  elements,  the 
oxidizing  effects  predominate.  Even  other  compounds  containing 
oxygen  will  give  exactly  the  same  results.  Hence  this  cannot  be  re- 
garded as  a  property  of  water  itself. 

Water  a  Stable  Compound.  —  Whether  the  substance  is  relatively 
stable  or  unstable  is,  in  the  case  of  a  compound,  the  first  chemical 
property  to  be  given.  Usually  the  specification  is  in  terms  of  the 
temperature  required  to  decompose  it.  Thus,  potassium  chlorate  gives 
off  oxygen  at  a  low  red  heat.  Water,  even  at  2500°,  is  only  1.8  per  cent 
decomposed,  and  reunion  occurs  when  the  temperature  is  lowered. 

Union  of  Water  with  Oxides.  —  When  sodium  combines  with 
oxygen  under  certain  conditions  we  obtain  sodium  oxide  (JSTa^O).  The 
product  unites  violently  with  water  to  form  sodium  hydroxide  : 


H2O  ->  2NaOH. 

The  slaking  of  quicklime  is  a  more  familiar  action  of  the  same  kind  : 
CaO  +  H20  -+  Ca(OH)2. 

No  other  products  are  formed.  The  clouds  of  steam  produced  in  the 
second  instance  are  due  to  evaporation  of  a  part  of  the  water  by  the 
heat  produced  in  the  formation  of  calcium  hydroxide.  The  aqueous 
solutions  of  these  two  products  have  a  soapy  feeling,  and  turn  red  lit- 
mus blue,  and  the  substances  therefore  belong  to  the  class  of  alkalies 
(p.  98)  or  bases.  Very  many  hydroxides  which  are  of  the  same  na- 
ture, for  example  ferric  hydroxide  (Fe(OH)8)  and  tin  hydroxide 
(Sn(OH)2),  are  formed  so  slowly  by  direct  union  of  the  oxide  and  water 
that  they  are  always  prepared  in  other  ways. 

Some  oxides,  although  they  unite  with  water,  give  products  of  an 
entirely  different  character.  Phosphoric  anhydride  and  sulphur  dioxide 
(p.  71)  are  of  this  class  and  yield  acids. 

These  two  classes  of  final  products  are  so  different  that  we  make 
the  distinction  the  basis  of  classification  of  the  elements  present  in  the 
original  oxides.  The  elements,  like  sodium  and  iron,  whose  oxides 
give  bases,  are  called  metals  ;  those,  like  phosphorus,  whose  oxides  give 
acids,  are  called  non-metals.  The  distinguishing  words  are  selected 
because  the  division  corresponds,  in  a  general  way  at  least,  with  the 
separation  into  two  sets  to  which  merely  physical  examination  of  the 
elementary  substances  would  lead. 


120  INORGANIC  CHEMISTRY 

Formerly  the  hydroxides  of  metals  were  termed  "  hydrates,"  and  the 
word  is  still  used  familiarly  by  chemists  in  a  few  cases,  such  as  potas- 
sium "  hydrate  "  (KOH)  and  sodium  «  hydrate  "  (NaOH).  These  sub- 
stances, however,  have  nothing  in  common  with  the  compounds  properly 
known  as  hydrates  whose  nature  is  discussed  in  the  next  section. 

Hydrates.  —  Many  substances  when  dissolved  in  water  and  re- 
covered by  spontaneous  evaporation  of  the  solvent  are  found  to  have 
entered  into  combination  with  the  liquid.  The  products,  which  are 
solids,  are  called  hydrates.  These  compounds  show  definite  chemical 
composition  expressible  in  terms  of  chemical  unit  weights  of  the  con- 
stituents. Often  much  heat  is  given  out  in  their 
formation.  Thus,  in  the  case  of  washing  soda, 
the  decahydrate  of  sodium  carbonate  (Na^CO,,, 
10H20),the  heat  of  the  union  (p.  76)  is  8800  cal. 
The  hydrates  have  physical  properties  entirely 
different  from  those  of  their  components.  Thus, 
cupric  sulphate,  often  called  anhydrous  cupric 
sulphate  to  distinguish  it  from  the  compound 
FIG  41  with  water,  is  a  white  substance  crystallizing 

in   shining,  colorless,  needle-like  prisms.     The 

pentahydrate  (blue-stone  or  blue  vitriol)  is  blue  in  color,  and  forms 
larger  but  much  less  symmetrical  (asymmetric  or  triclinic)  crystals 
(Fig.  41): 

CuS04  +  5H20-»  CuS04, 


When  heated,  the  hydrates,  as  a  rule,  lose  none  of  the  constituents  of 
the  original  compound,  but  only  those  of  the  water.  They  do  so,  in 
general,  rather  easily.  Hence,  to  avoid  the  disguise  of  the  fundamental 
substance  which  would  occur  if  we  wrote,  for  example,  H10CuS09,  we 
segregate  (as  above)  the  elements  of  the  water  in  the  formula.  The 
aqueous  solutions  made  from  the  anhydrous  substances  and  from  the  hy- 
drates have  identical  physical  and  chemical  properties.  Hence  the 
cheaper  of  the  two  forms  is  generally  purchased,  and  many  of  the 
chemicals  used  in  laboratories  are  in  the  form  of  hydrates. 

Some  of  these  hydrates  decompose  very  readily.  The  decahydrate 
of  sodium  sulphate,  Na2S04,  10H20  (Glauber's  salt),  gives  up  all  the  water 
it  contains  when  simply  kept  in  an  open  vessel.  At  100°  blue  vitriol 
very  quickly  loses  4H20  and  the  rest  of  the  water  more  slowly.  The 
last  equation  might  therefore  have  been  written  to  show  a  reversible 


WATER  121 

action  (p.  64).  A  decomposition  which  proceeds  at  high  temperatures, 
while  at  lower  temperatures  recombination  of  the  constituents  can  take 
place,  as  with  these  hydrates,  is  called  a  dissociation.  The  decomposi- 
tion of  potassium  chlorate  (p.  64)  is  not  a  dissociation  because  it  is  not 
reversible ;  oxygen  will  not  under  any  circumstances  reunite  with  potas- 
sium chloride. 

The  condition  which  controls  such  actions  is  not,  however,  the  tem- 
perature alone.  When  Glauber's  salt  is  kept  in  a  closed  bottle,  a  very 
little  of  it  loses  water,  and  then  the  decomposition  ceases.  When  the 
bottle  is  left  open,  the  dissociation  proceeds  until  no  decahydrate  re- 
mains. The  cause  of  this  we  discover  when  a  crystal  of  the  hydrate 
is  placed  above  mercury,  like  the  ice  or  water  in  Fig.  39  (p.  116).  It 
shows  a  definite  aqueous  tension.  At  9°  the  value  of  this  is  5.5  mm. 
As  its  temperature  is  raised,  the  tension  increases.  When  the  tem- 
perature is  lowered,  on  the  other  hand,  the  tension  diminishes,  the 
mercury  rises,  and  more  of  the  water  enters  into  combination  again. 
Different  hydrates  show  different  aqueous  tensions  at  the  same  tempera- 
ture. For  example,  at  30°,  water  itself  is  31.5  mm.,  strontium  chloride 
(SrCljj,  6H20)  11.5  mm.,  cupric  sulphate  (CuSO4, 5H20)  12.5  mm.,  barium 
chloride  (BaCLj,  2H20)  4  mm. 

It  appears,  therefore,  that  the  water  in  these  compounds  evaporates 
as  do  ordinary  waters.  Those  which,  like  washing  soda,  have  a  vapor 
tension  approaching  that  of  water  itself,  lose  their  water  at  ordinary 
temperatures  at  a  rapid  pace.  In  this  connection  we  have  to  remember 
that  atmospheric  air  is  usually  less  than  two-thirds  saturated  with  water 
vapor,  and  the  partial  pressure  of  this  vapor  opposes  the  dissociation. 
Thus  at  9°,  the  vapor  tension  of  water  being  8.6  mm.,  the  average  vapor 
pressure  of  water  in  the  atmosphere  will  be  about  5  mm.  Any  hydrate 
with  a  greater  aqueous  tension  than  5  mm.,  at  9°,  such  as  Glauber's 
salt,  will  therefore  decompose  spontaneously  in  an  open  vessel.  But 
those  with  a  lower  vapor  tension,  such  as  the  pentahydrate  of  cupric 
sulphate  with  a  tension  of  2  mm.  at  9°,  will  not  do  so  (see  p.  135). 

The  behavior  of  hydrates  does  not  indicate,  as  might  seem  at  first 
sight  to  be  the  case,  that  the  water  is  contained  in  them  in  some  way 
in  the  free  state.  The  fact  is  that  the  above  statements,  with  corre- 
sponding changes  in  the  wording,  might  be  made  of  all  dissociations  in 
chemistry.  Oxides  give  a  different  pressure  of  oxygen  at  each  tem- 
perature, carbonates  of  carbon  dioxide,  and  so  forth. 

The  measurement  of  the  vapor  tension  of  hydrates  gives  definite  information 
in  regard  to  whether  there  are  other  hydrates,  say  of  cupric  sulphate,  with  less 


122  INORGANIC   CHEMISTRY 

than  the  normal  number  of  formula-weights  of  water.  If  there  were  only  two 
substances,  CuS04  and  CuS04,  5H2O,  with  no  compound  of  intermediate  composi- 
tion, then  a  partially  decomposed  specimen  would  be  made  up  partly  of  the  one 
substance  and  partly  of  the  other.  But  if  there  were  an  intermediate  compound, 
say  CuS04,  3H2O,  then  desiccating  a  specimen  of  the  pentahydrate  would  give 
nothing  but  mixtures  of  CuSO4,  3H2O  and  CuS04,  5H20  until  all  the  latter  was 
decomposed.  Then,  and  only  then,  the  trihydrate  would  begin  to  lose  water.  Now 
the  trihydrate,  being  a  definite  and  different  substance,  would  have  a  vapor  tension 
of  its  own,  and  experimental  study  would  show  its  presence. 

Experiment  shows  that  there  really  are  several  hydrated  cupric  sulphates. 
The  pentahydrate,  at  50°,  has  a  vapor  tension  of  47  mm.,  and  this  vapor  tension  is 
observed  so  long  as  any  pentahydrate  remains  to  be  decomposed.  As  soon  as  the 
proportion  of  water  goes  down  to  CuS04,  3H2O,  the  vapor  tension  suddenly  drops 
to  30  mm.  As  the  desiccation  continues,  this  tension  is  maintained  until  the  com- 
position has  reached  CuS04,  H2O.  At  this  point  the  vapor  pressure  falls  to  that 
of  the  monohydrate,  4.5  mm.,  and  remains  at  this  value  until  all  the  rest  of  the 
water  has  been  removed.  Had  there  been  no  intermediate  compound  with  3H20 
the  tension  would  have  dropped  at  once  from  47  mm.  to  4.5  mm.  If,  conversely,  we 
try  to  combine  water  as  vapor  with  anhydrous  cupric  sulphate,  at  50°,  a  vapor 
pressure  of  at  least  4.5  mm.  is  required  to  cause  union  to  take  place.  The  union 
stops  when  one  formula-weight  of  water  has  undergone  combination.  To  intro- 
duce more,  the  concentration  of  the  water  vapor  must  be  increased  to  nearly  seven 
times  its  first  value,  namely,  to  30  mm.  pressure.  This  enforces  combination  up  to 
CuS04,  3H2O.  For  further  hydration,  a  still  higher  pressure  of  water  vapor  is 
needed  (47  mm.),  and  the  absorption  ceases  when  CuS04,  5H2O  has  been  formed. 

There  are  thus  three  distinct  reversible  actions  which  succeed  one  another  as 
the  hydration  proceeds : 

CuS04  +  H2O  T=±  CuS04,  H20 
CuS04,  H2O  +  2H2O  ^±  CuSO4,  3H2O 
CuS04,  3H20  +  2H2O  <=»  CuSO4,  5H2O. 

The  first  represents  a  greater  affinity  than  the  second,  and  the  second  than  the 
third. 

The  graphic  representation  of  these  facts  (Fig.  42)  will  make  the  behavior  of  the 
compounds  clearer.  The  proportion  of  water  combined  with  one  formula-weight 
of  cupric  sulphate  is  laid  off  along  the  horizontal  axis.  The  pressures  at  which  it 
enters  or  leaves  the  compounds  at  50°  are  the  ordinates.  As  far  as  1H20  the  pressure 
is  constant  (4.5  mm.).  Beyond  that  point  and  up  to  3H2O  it  is  constant  but  much 
higher.  Between  3H2O  and  5H2O  it  is  constant  again  but  higher  still. 

The  tension  of  free  water  at  the  same  temperature  is  92  mm.  It  is  constant 
irrespective  of  the  amount  of  water,  and  would  therefore  be  on  a  single  continuous 
line  parallel  to  the  horizontal  axis  and  twice  as  high  above  it  as  the  uppermost  one 
in  the  diagram.  If,  at  50°,  a  vessel  of  water  were  put  under  a  bell  jar  alongside  of 
anhydrous  cupric  sulphate,  its  vapor  would  be  more  than  sufficiently  concentrated 
fully  to  hydrate  the  compound.  Again,  while  4.5  mm.  pressure  of  water  vapor 
will  cause  water  to  combine  with  anhydrous  cupric  sulphate  at  50°,  a  pressure  of 
92  mm.  will  be  required  to  liquefy  the  water  vapor  at  the  same  temperature. 

The  above  discussion  shows  that  the  last  formula- weight  of  water,  in  hydrates 


WATER 


128 


50 

47— 


40 


s 


30 


20 


10 


4.5 
0 


which  dissociate  by  stages,  is  not  different  in  kind  from  the  others.     It  differs  only 

in   the  degree  of    tenacity  with  which  it  is  held.     It  is  therefore  unnecessary, 

merely  on  this  account,  to  dignify  it  by  the  separate  name  of  water  of  constitu- 
tion, as  has  been  done  by  some  chemists. 

Water  of  hydration  is  frequently  called  water  of  crystallization,  on  account 

of  the  fact  that  when  water  is  driven  off  by  heating,  the  substance  usually  crumbles 

to  pieces  (effloresces) .    The  term  is  not  particularly  appropriate  for  several  reasons. 

It  suggests  that  water  and  crys- 
tallization are  related  in  some 

way,  which  is  not    the    case. 

Sulphur,      galena,      potassium 

chlorate,  and  thousands  of  other 

crystallized  substances,  do  not 

contain  the  elements  of  water. 

Nor  do   the  substances  which 

combine    with    water    remain 

amorphous     in     its     absence. 

They   all   crystallize  from   the 

molten  condition  or  from  some 

non-aqueous  solvent,  although, 

as  substances  different  from  the 

hydrates,  their  crystalline  form 

is  different.      Iceland  spar,  or 

any  other  crystallized  carbonate 

which  can   be  decomposed  by 

heating,  becomes    opaque  and 

porous  or  falls  to  powder  when  the  carbon  dioxide  is  driven  out.     But  it  has 

not  occurred  to  any  one  to  call  this  carbon  dioxide  of  crystallization!    The  fact  is 

that  all  pure  chemical  substances,  in  solid  form,  when  in  a  stable  physical  condi- 
tion, are  crystalline.     Amorphous  substances  are  always  supercooled  liquids. 

The  term  arose  from  a  misconception,  and,  when  used,  always  succeeds  in 
transmitting  the  misconception  along  with  the  name.  The  ease  with  which  some 
of  the  hydrates  decomposed  suggested  the  idea  that  they  contained  water  as  a  dis- 
crete substance.  There  is  no  more  justification  for  this  idea,  however,  than  for  the 
notion  that  carbonates  contain  ready-made  carbon  dioxide.  The  hydrates  contain 
the  elements  of  water  just  as  sugar  and  alcohol  do,  and  there  is  no  evidence  that 
they  "contain  water"  in  any  other  sense  than  that  in  which  the  phrase  might  be 
used  of  these  organic  bodies. 

In  consequence  of  their  decomposition  into  and  formation  from  substances 
capable  of  separate  existence,  the  hydrates  are  classed  with  molecular  compounds 
(q.v. ).  The  behavior  of  the  compounds  of  salts  with  ammonia  (like  2AgCl,  SNH^, 
with  nitric  oxide,  and  with  each  other  (double  salts),  is  quite  similar. 

Composition  of  Water.  — The  measurement  of  the  proportions 
by  weight  and  volume  in  which  hydrogen  and  oxygen  combine  to  form 
water  has  been  the  subject  of  a  larger  number  of  elaborate  investiga- 
tions than  any  other  single  problem  of  this  kind.  The  difficulty  in 


1234 
Formula-weights  of  water 

FIG.  42. 


124 


INOKGANIC   CHEMISTRY 


making  the  former  measurement  arises  from  the  fact  that  both  constit- 
uents are  gases,  and  therefore  difficult  to  weigh. 

The  determination  of  this  proportion  which,  until  very  recently,  held  its  place 
in  all  chemical  works,  was  that  made  by  the  French  chemist,  Dumas.  His  experi- 
ments gave  the  ratio  of  hydrogen  to  oxygen  2  : 15.96. 

It  was  not  until  1887  that  Reiser  obtained  a  figure  for  the  oxygen  appreciably 
smaller  than  this,  and  soon  determinations  by  other  observers  showed  that  Dumas' 
proportion  was  probably  too  large.  The  investigation  which 
finally  settled  this  question  was  that  of  Edward  Morley.  The 
most  striking  of  his  experiments  consisted  in  a  series  of  syn- 
theses of  water,  in  which  he  weighed  the  hydrogen  as  well  as 
the  oxygen,  and  afterwards  weighed  the  water  produced  from 
them.  The  hydrogen  was  confined  by  absorption  in  palladium 
(p.  107),  and  could  thus  be  contained  in  large  quantity  in  a 
small,  elongated  bulb.  During  the  progress  of  the  experiment 
it  was  driven  out  by  a  suitable  heating  arrangement.  The 
oxygen  was  contained  in  large  globes  holding  15-20  liters.  The 
losses  in  weight  of  the  palladium  tube  and  of  the  globes  gave 
the  hydrogen  and  oxygen  consumed.  The  manipulator  in 
which  the  gases  were  combined  and  the  water  collected  is 
represented  in  Fig.  43.  The  gases  entered  through  two  small 
tubes  marked  A.  Just  above  them,  between  two  platinum 
wires,  a  discharge  of  electricity  started  the  union  and  when 
necessary  maintained  it.  The  vessel  was  first  filled  by  admit- 
ting oxygen,  and  the  hydrogen  was  burned  at  the  mouth  of 
the  tube  from  which  it  issued.  This  part  of  the  apparatus 
was  immersed  in  a  vessel  of  water  with  transparent  walls 
through  which  the  union  could  be  watched,  and  the  steam 
formed  was  condensed  and  collected  in  the  bottom  of  the  ves- 
sel. The  vacuum  thus  produced  enabled  the  oxygen  continually 
to  flow  into  the  manipulator  from  the  globes.  In  this  way 
forty-two  liters  of  hydrogen  and  twenty -one  liters  of  oxygen 
could  be  combined  in  about  an  hour  and  a  half. 

At  the  end  of  the  experiment  this  part  of  the  apparatus 
was  disconnected  and  placed  in  a  freezing  mixture  which  con- 
verted the  water  into  ice  and  practically  condensed  the  whole 
of  its  vapor.  The  uncombined  gas  in  the  apparatus  was  withdrawn  and  its 
nature  and  quantity  determined.  The  increase  in  weight  of  the  manipulator 
gave  the  quantity  of  water  formed.  The  success  of  each  experiment  could  be 
tested  by  comparing  the  sum  of  the  weights  of  oxygen  and  hydrogen  with  that 
of  the  water  obtained  from  them.  The  manipulation  was  so  skilful,  and  the 
various  corrections  used  were  so  adequate,  that  this  difference  was  almost 
negligible.  The  ratio  of  hydrogen  to  oxygen  in  water  in  this  series  of  experi- 
ments was  2  : 15.879,  a  result  which  agreed  with  the  other  methods  of  determin- 
ing the  same  ratio  which  Morley  used.  It  agrees  also  exactly  with  the  average 
of  the  numbers  obtained  by  several  other  observers. 

The  most  probable  value  of  the  ratio  by  weight,  taking  his  own 


FIG.  43. 


WATER 


125 


and  other  trustworthy  measurements  into  account,  is  given  by  Morley 
as  2 : 15.879  or  2.015 : 16.  The  proportion  by  volume  is  2.0027 
volumes  of  hydrogen  to  1  volume  of  oxygen. 

That  the  proportion  by  volume  is  very  close  to  2  : 1  may  easily  be 
shown.  We  may  use  a  U-shaped  tube  closed  at  one  end  by  a  stopcock 
and  graduated  (Fig.  44).  At  first,  the  left  limb  of  the  tube,  called  a 
eudiometer,  is  filled  with  mercury.  One  of  the  gases  is  admitted  so  as  to 
fill  a  portion  of  the  tube  and,  the  levels  having  been 
equalized  (cf.  p.  82),  the  volume  of  the  gas  is  read. 
Then  some  of  the  other  gas  is  introduced  and  the 
leveling  and  reading  repeated.  Let  us  suppose 
that  15  c.c.  of  hydrogen  and  10  c.c.  of  oxygen 
have  thus  been  taken.  The  right  limb  is  then 
filled  with  mercury  and  closed  firmly  with  the 
thumb.  A  spark  from  an  induction  coil  passing 
between  the  two  short  platinum  wires  near  the  top 
of  the  tube  explodes  the  mixture.  The  steam  pro- 
duced by  the  union  condenses  almost  immediately 
and  occupies  practically  no  volume  worth  con- 
sidering. When  the  thumb  is  removed,  the  mer- 
cury rises  on  the  left  and  fills  up  the  space  left 
by  the  disappearance  of  part  of  the  gases.  Unless 
the  proportion  taken  happens  to  have  been  exact, 
some  of  one  or  other  of  the  gases  will  remain. 
Its  volume  is  measured  by  equalizing  the  levels 
and  reading  as  before.  In  the  case  we  have  im- 
agined, the  residual  gas  is  oxygen,  and  there  are 
almost  exactly  2.5  c.c.  of  it.  It  is  evident,  there- 
fore, that  15  c.c.  of  hydrogen  united  with  7.5  c.c.  of  oxygen ;  in  other 
words,  the  proportion  by  volume  is  2  :  1. 

Gay-Lussac's  Law  of  Combining  Volumes. —  The  almost 
mathematical  exactness  with  which  small  integers  express  this  propor- 
tion is  not  a  mere  coincidence.  Whenever  gases  unite,  or  gaseous 
products  are  formed,  the  proportion  by  volume  (measured  at  the 
same  temperature  and  pressure)  of  all  the  gaseous  bodies  concerned 
can  be  represented  very  accurately  by  ratios  of  small  integers.  This 
is  called  Gay-Lussac's  law  of  combining  volumes  (1808).  Thus,  when 
the  above  experiment  is  carried  out  at  100°,  in  order  that  the  product, 
water,  may  be  gaseous  also,  it  is  found  that  the  three  volumes  of  the 


FIG.  44. 


126 


INORGANIC   CHEMISTRY 


constituents  give  almost  exactly  two  volumes  of  steam.  For  example, 
15  c.c.  of  hydrogen  and  7.5  c.c  of  oxygen  give  15  c.c.  of  steam. 
Of  course  the  hydrogen,  oxygen,  and  steam  must  be  measured  at 

the  same  pressure,  and  the  temperature 
must  remain  constant  (100°)  during  the 
experiment.  Proper  manipulation  secures 
the  former,  and  a  jacket  filled  with  steam 
(Fig.  45)  the  latter  condition.  Strips  of 
paper,  1,  2,  and  3,  are  pasted  on  the  jacket 
in  such  a  way  that  equal  lengths  of  the 
eudiometer,  in  this  case  a  straight  one, 
are  laid  off.  The  three  divisions  having 
been  filled  with  a  mixture  of  hydrogen 
and  oxygen  in  the  proper  proportions,  the 
gas,  after  the  explosion,  shrinks  so  as  to 
occupy,  at  the  same  pressure,  only  two 
of  them. 

From  this  universal  truth  in  regard 
to  the  combination  of  gases,  we  draw  the 
important  inference  that  the  chemical 
unit-weights  of  simple  substances,  and  the 
formula-weights  of  compounds,  in  the 
gaseous  condition,  occupy  at  the  same 
temperature  and  pressure  volumes  which 
are  equal  or  stand  to  one  another  in  the 
ratio  of  small  integers  (see  Molar  weights). 

FIG.  45. 

The  chemical  behavior  of  the  other  compound 
of  hydrogen  and  oxygen,  hydrogen  peroxide 

(<?.».),  is  difficult  to  comprehend  until  further  experience  has  been  gained. 
Then,  too,  its  formula  (H2O2)  cannot  be  justified  until  the  means  of  deter- 
mining molar  weights  in  solution  have  been  discussed.  In  view  of  these  facts, 
and  because  it  furnishes  the  simplest  illustration  of  the  meaning  of  a  structural 
formula,  a  conception  which  would  be  quite  out  of  place  at  this  stage,  it  will  be 
taken  up  later. 


Exercises.  —  1.  Name  some  other  transitions  from  one  physical 
state  to  another,  which  are  familiar  (p.  115). 

2.  What  evidence  is  there  in  the  common  behavior  of  ether, 
alcohol,  and  chloroform  tending  to  show  that  these  liquids  have  high 
vapor  tensions  ? 


WATER  127 

3.  If  the  pressure  of  the  steam  in  a  boiler  is  ten  atmospheres,  at  what 
temperature  is  the  water  boiling  (p.  116)  ? 

4.  How  many  grams  of  water  could  be  heated  from  20°  to  100°  by 
the  heat  required  to  melt  1  kgm.  of  ice  at  0°  ? 

5.  What  do  you  infer  from  the  fact  that  alum  and  washing  soda 
lose  their  water  of  crystallization  when  left  in  open  vessels,  while 
gypsum  does  not  (p.  121)  ? 

6.  Which  facts  show  conclusively  that  hydrates  are  true  chemical 
compounds?     Is   there  any  fact  which   throws    doubt   on    this    con- 
clusion ? 

7.  In  what  way  does  a  hydrate  differ  from  (a)  a  solution,  (b)  an 
hydroxide  ? 

8.  Should  you  expect  to  find  any  difference,  in  respect  to  chemical 
activity,  between  the  three  forms  of   water  (ice,  water,  and  steam)  ? 
If  so,  arrange  them  in  the  order  of  probable  increasing  activity  (pp. 
26-28).     Have  we  had  any  experimental  confirmation,  or  the  reverse, 
of  this  conclusion  ? 

9.  Which  contains  more  chemical  energy,  and  is  therefore  more 
active,  the  anhydrous  substance,  or  the  corresponding  hydrate  ? 


CHAPTER   IX 

THE   KINETIC-MOLECULAR   HYPOTHESIS 

As  soon  as  we  have  constructed  a  law  (p.  7)  we  desire  immediately 
to  find  out  the  basis  of  the  constant  mode  of  behavior  it  epitomizes. 
If  no  explanation,  that  is,  more  detailed  description,  is  forthcoming 
as  the  result  of  closer  observation,  we  proceed  to  imagine  one  (p.  10). 
This  always  takes  a  mechanical  form,  often  crude  at  first,  and  later 
undergoing  refinement.  Thus,  at  first,  the  phenomena  of  light  were 
explained  by  the  conception  of  clouds  of  fine  corpuscles  emanating 
from  the  luminous  body.  The  chances  of  hitting  upon  an  objective 
reality  by  guess-work  like  this  is  obviously  remote.  Whether  such 
particles  did  really  fly  about  was  not  the  main  question,  however. 
Their  value  lay  in  the  fact  that  they  could  be  pictured  concretely  and 
gave  a  basis  for  further  thought  and  perhaps  suggestions  for  new 
experiments.  Such  a  structure  of  the  imagination  is  called  an  hypoth- 
esis. If  it  furnishes  an  explanation  of  more  than  one  law,  so  much 
the  better. 

The  Molecular  Hypothesis.  —  The  only  mechanical  basis  we  can 
imagine  to  account  for  the  physical  properties  of  matter  is  a  discontin- 
uous structure  of  some  description.  The  fact  that  all  kinds  of  matter 
can  be  compressed  (gases  to  an  enormous  extent,  solids  and  liquids  to  a 
measurable  extent)  may  be  explained,  either  by  a  diminution  in  the 
volume  of  the  material  itself,  or  by  the  closer  packing  together  of  the  par- 
ticles into  which  this  material  is  divided.  It  is  evident  that  the  latter 
is  much  more  in  harmony  with  our  experience.  Compression,  we 
imagine,  therefore,  does  not  diminish  the  actual  volume  occupied  by 
matter,  but  crowds  the  particles  closer  together  and  diminishes  the 
space  between  them.  The  same  hypothesis  will  furnish  a  concrete 
description  of  how  one  kind  of  matter  will  frequently  absorb  a  large 
quantity  of  another.  Thus  we  picture  the  hydrogen  gas  taken  up  by 
iron  and  other  metals  as  being  packed  away  in  the  spaces  between  the 
particles  of  the  metal.  So  also,  in  solution,,  the  volume  of  the  liquid 
does  not  usually  increase  by  an  amount  equal  to  that  of  the  substances 

128 


THE    KINETIC-MOLECULAR   HYPOTHESIS  129 

dissolved.  Hence  we  imagine  the  particles  as  possibly  incompressible 
and  the  interstices  between  them  as  furnishing  a  part  of  the  accommo- 
dation for  the  foreign  material.  A  particle  is  a  fragment  of  matter 
which  can  be  seen  and  measured  directly,  and  handled  separately. 
Since  the  particles  of  this  hypothesis  have  none  of  these  qualities,  we 
distinguish  them  by  the  name  molecules.  Molecules  are  the  imaginary 
units  of  -which  bodies  are  aggregates. 

Since  the  behavior  of  gases  has  been  considered  most  fully  in  the 
preceding  chapters,  and  is  in  any  case  capable  of  more  exact  and 
simple  description  than  that  of  solids  or  liquids,  we  shall  apply  this 
hypothesis  first  to  them. 

Kinetic-Molecular  Hypothesis  Applied  to   Gases*  —  Let   us 

first  build  up  our  hypothesis  to  fit  the  qualitative  properties  of  gases. 
The  most  remarkable  thing  about  a  gas,  considering  the  looseness  with 
which  its  material  is  packed,  is  the  total  absence,  in  it  of  any  tendency 
to  settling  or  subsidence.  Since  the  molecules  cannot  be  at  rest  upon 
one  another,  as  the  great  compressibility  shows,  we  are  driven  to 
suppose  that  they  are  widely  separated  from  one  another,  and  that 
they  occupy  the  space  by  constantly  moving  about  in  all  directions. 
But  a  moving  aggregate  of  particles  which  does  not  even  finally  settle 
must  be  in  perpetual  motion.  We  must,  therefore,  imagine  the  mole- 
cules to  be  wholly  unlike  particles  of  matter  in  having  perfect  elasticity, 
in  consequence  of  which  they  undergo  no  loss  of  energy  after  a 
collision.  They  must  continually  strike  the  walls  of  the  vessel  and 
one  another  and  rebound,  yet  without  loss  of  motion.  The  diffusibility 
of  gases  and  their  mutual  permeability  require  no  additional  assump- 
tions. The  fact  that  each  gas  is  homogeneous,  efforts  to  sift  out  lighter 
or  heavier  samples  having  failed,  requires  the  supposition  that  all  the 
molecules  of  a  pure  gas  are  closely  alike. 

Passing  now  to  Boyle's  law  (p.  84),  the  thing  to  be  accounted  for 
is  that  when  a  sample  of  a  gas  diminishes  in  volume,  its  pressure 
increases  in  the  same  proportion.  Let  the  diagram  (Fig.  46)  represent 
a  cylinder  with  a  movable  piston,  upon  which  weights  may  be  placed 
to  resist  the  pressure.  Now  the  pressure  exercised  by  the  gas  cannot 
be  like  the  pressure  of  the  hand  upon  a  table,  since  we  have  just 
assumed  that  the  particles  are  not  even  approximately  at  rest,  and  the 
spaces  between  them  are  enormous  compared  with  the  size  of  the 
molecules  themselves.  The  gaseous  pressure  must  therefore  be  attri- 
buted to  the  colossal  hailstorm  which  their  innumerable  impacts  upon 


130  INORGANIC   CHEMISTRY 

the  piston  produce.  If  this  is  the  case,  the  compressing  of  a  gas  must 
consist  simply  in  moving  the  partition  downwards  so  that  the  particles 
as  they  fly  about  are  gradually  restricted  to  a  smaller  and  smaller 
space.  Their  paths  become  on  an  average  shorter  and  shorter.  Their 
impacts  upon  the  wall  become  more  and  more  frequent.  So  the 
pressure  which  this  occasions  becomes  greater  and  greater,  and  is  pro- 
portional to  the  degree  of  crowding  of  the  molecules. 
There  are  two  other  points  which  must  be  added.  When 
we  diminish  the  volume  to  one-half,  we  find  from  experience 
that  the  pressure  becomes  exactly,  or  almost  exactly,  twice 
as  great.  This  must  mean  that  although  the  particles  are 

T  becoming  crowded  they  do  not  interfere  with  one  another's 

motion,  excepting  of  course  where  actual  collision  causes 
a  rebound.  Only  in  the  absence  of  interference  would 
doubling  the  number  of  particles  per  unit  of  volume  give 
exactly  double  the  number  of  impacts  on  the  walls.  Hence 


FIG.  46.  the  particles  must  have  practically  no  tendency  to  cohesion. 
Again,  the  molecules  must  move  in  straight  lines,  because, 
if  they  moved  in  orbits  of  some  kind,  many  of  the  orbits  would  not  be 
intersected  by  the  wall  of  the  vessel  until  great  reduction  in  the 
volume  had  taken  place,  and  thus,  as  the  volume  diminished,  the  fre- 
quency of  the  impacts,  and  therefore  the  pressure,  would  increase 
faster  than  the  concentration. 

Boyle's  law  therefore  adds  four  more  conceptions  to  our  molecular 
hypothesis,  namely,  that  the  impacts  of  the  particles  produce  the  press- 
ure, that  the  crowding  of  the  molecules  represents  the  concentration 
(p.  80),  and  that  the  particles  move  in  straight  lines  and  show  almost 
no  cohesion,  since  pressure  and  concentration  are  very  closely  propor- 
tional to  one  another. 

It  will  be  seen,  on  consideration,  that  if  the  molecules  are  assumed  to  repel 
one  another,  they  would  do  so  more  violently  the  more  closely  they  were  packed 
together.  This  assumption  would  therefore  suit  the  case  of  a  gaseous  body  in 
which  the  pressures  increased  according  to  some  power  of  the  concentration  other 
than  the  first,  and  therefore  much  more  rapidly  than  in  known  gases.  In  spite  of 
its  inapplicability,  this  notion  is  supposed  by  many  people  to  be  part  of  the 
kinetic  hypothesis. 

Charles'  law  (p.  87),  that  a  gas  receives  equal  increments  in  volume 
or  pressure  for  equal  elevations  in  temperature,  requires  but  one  addi- 
tion to  the  .hypothesis.  Concretely,  if  our  specimen  of  gas  (Fig.  46)  is 
at  0°,  and  we  permit  its  pressure  to  remain  constant  by  leaving  the 


THE   KINETIC-MOLECULAR   HYPOTHESIS  131 

same  weight  on  the  piston,  then  when  the  temperature  of  the  gas  is 
raised  to  1°,  the  volume  will  gain  ^^  of  the  original  volume.  If,  on 
the  other  hand,  we  restrict  the  gas  to  the  original  volume,  the  pressure 
will  evidently  increase,  and  the  augmentation  will  be  ^\^  of  the 
original  pressure.  Now,  how  can  we  account  for  an  increase  in 
pressure  as  the  result  of  heating  a  mass  of  rapidly  moving  molecules  ? 
The  action  of  a  particle  colliding  with  a  surface  is  measured  in  physics 
in  terms  of  its  mass  and  its  velocity.  It  is  evident  that  heating  a  cloud 
of  molecules  would  not  increase  the  mass  of  each,  and  it  must  there- 
fore increase  the  velocity  of  each  since  the  kinetic  energy  of  all 
becomes  greater.  This  conclusion  is  in  harmony  with  our  experience 
that  violently  rubbing  a  solid  raises  its  temperature,  and  such  a  mode 
of  treatment  might  plausibly  be  supposed  to  communicate  motion  to 
the  minute  parts  of  the  body. 

The  fact  that  the  combining  volumes  of  gaseous  substances  are  equal, 
or  stand  to  one  another  in  the  ratio  of  small  whole  numbers  (cf.  Gay- 
Lussac's  law,  pp.  125-126),  suggests  two  ideas :  First,  that  chemical 
combination,  considered  in  detail,  and  arranged  to  harmonize  with  this 
hypothesis,  would  involve  unions  of  a  few  particles  of  more  than  one 
kind  to  form  composite  molecules.*  And,  second,  that  a  simple  inte- 
gral relation  must  be  assumed  to  exist  between  the  numbers  of  mole- 
cules in  equal  volumes  of  different  gases,  at  the  same  temperature  and 
pressure.  Avogadro  (1811),  the  professor  of  physics  in  Turin,  put 
forward  the  hypothesis  that  these  numbers  might  be  equal.  A  more 
strict  study  of  the  assumptions  we  have  been  making,  and  of  some  addi- 
tional facts,  has  since  shown  that  no  other  conjecture  than  Avogadro's 
would  be  consistent  with  them.  Thus  it  now  bears  the  relation  of  a 
logical  deduction  from  the  kinetic-molecular  hypothesis  and  the 
properties  of  gases,  and  is  known  as  Avogadro's  hypothesis.  It  may 
also  be  put  in  the  form  :  At  the  same  temperature  and  pressure,  the 
molecular  concentration  (cf.  p.  80)  of  all  kinds  of  gases  has  the  same 
value. 

The  law  of  diffusion  (p.  108)  harmonizes  with  the  kinetic-molecular 
hypothesis  without  further  modification  of  the  latter.  The  strict 
deduction  of  this  law,  as  well  as  of  the  preceding  ones,  from  our  series 
of  assumptions,  will  be  found  in  any  work  on  physical  chemistry. 

Finally,  we  have  referred  (p.  85)  to  the  fact  that  at  low  pressures 

*  This  is  essentially  the  idea  used  by  Dalton,  before  Gay-Lussac's  law  was 
known,  however,  for  the  explanation  of  the  laws  of  chemical  combination.  He 
called  it  the  atomic  hypothesis  (g.«.). 


132  INORGANIC    CHEMISTRY 

the  concentration  increases  ?nore,  and  at  high  pressures  much  less  than 
Boyle's  law  indicates.  The  former  effect  is  brought  into  accord  with 
our  hypothesis  when  we  remember  that  the  matter  even  of  gases  can 
cohere,  as  is  shown  plainly  when  they  are  solidified.  The  tendency  of 
the  molecules  to  cohere  must  therefore  show  itself  in  the  gaseous  con- 
dition by  pulling  the  gas  together  and  producing  somewhat  greater 
concentration  than  is  strictly  consistent  with  the  value  of  the  pressure. 
Although  this  effect  of  cohesion  is  usually  insignificant,  the  modern 
method  of  liquefying  gases  (q.v.)  depends  upon  it  almost  entirely. 

The  abnormally  small  reductions  in  volume  which  occur  when  the 
volume  of  the  gas  has  already  been  greatly  reduced  remind  us  that,  ac- 
cording to  our  hypothesis,  it  is  only  the  space  between  the  molecules 
that  is  diminished  as  pressure  rises,  and  not  the  space  occupied  by  the 
molecules.  Hence,  when  the  molecules  have  become  so  crowded  to- 
gether that  this  irreducible  space  begins  to  form  an  appreciable  frac- 
tion of  the  whole,  a  doubling  of  the  pressure  will  diminish  to  one-half 
its  value  only  a  part  (the  vacant  part)  of  the  volume  the  gas  occupies. 

If  the  incompressible  space  occupied  by  the  molecules  is  called  6,  and  that  of  the 
whole  gas  i>,  then  the  amended  form  of  Boyle's  law  reads  p  (•»  — 6)  =  constant.  Sim- 
ilarly, if  the  cohesive  tendency  is  taken  into  account,  it  is  plain  that  its  effect  will 
be  numerically  greater  at  small  volumes,  although  not  so  easily  observed.  It  is  in 
fact  inversely  proportional  to  the  square  of  the  volume.  If  it  is  expressed  in  the 
same  units  as  the  pressure  by  a,  the  total  of  the  compressing  tendencies  becomes 

p  H — -  •     Hence  Boyle's  law,  for  constant  temperatures,  as  amended  by  Van  der 

Waals,  reads  fp+-J(t>  —  &)  =  constant,  a  formula  which  describes  the  actual 

behavior  of  most  gases  with  remarkable  accuracy.  Hydrogen  alone,  at  ordinary 
temperatures,  shows  no  excessive  compressibility  at  low  pressures.  The  cohesion 
(a)  is  so  slight  that  the  effect  of  the  constant  (6)  counterbalances  it  from  the  very 
first. 

We  may  summarize  the  facts  about  gases,  appearing  in  italics 
above,  with  the  corresponding  fictions,  in  heavy  type,  which  we  have 
added  one  by  one  in  manufacturing  our  hypothesis,  as  follows  : 


FACTS. 


HYPOTHESIS. 


Non-settling  .  . . 
Compressibility , 
Diff  usibility 
Permeability  . . , 


Homogeneity 


f  The  fictitious  particles  called  molecules  are, 
at  0°  and  760  mm. ,  at  great  average  distances 
from  one  another;  they  are  in  constant 
motion  and  have  perfect  elasticity. 

The    molecules    of    the    same    substance    are 
closely  alike. 


THE    KINETIC-MOLECULAR   HYPOTHESIS 


133 


FACTS. 


HYPOTHESIS. 


Relation  of  pressure  and  concen- 
tration (Boyle's  law). 


Relation  of  volume  (or  pressure) 
and  temperature  (Charles'  law). 


Relation  of  atomic  weights  and 
volumes  (Gay-Lussac's  law). 


Law  of  diffusion. 

Abnormal  incompressibility,  es- 
pecially at  high  pressures. 

Abnormal    compressibility,    es- 
pecially at  low  pressures. 


The  effect  of  pressure  is  produced  by  the  im- 
pacts of  the  molecules,  and  is  proportional  to 
the  degree  to  which  they  are  crowded  to- 
gether ;  the  molecules  move  in  straight  lines, 
and  have  almost  no  tendency  to  cohesion. 

A  rise  in  temperature  increases  the  velocity 
and  therefore  tho  kinetic  energy  of  the  mole- 
cules. 

Chemical  union  consists  in  fusion  of  different 
kinds  of  molecules  (Dalton's  hypothesis),  of 
which  there  are  equal  numbers  in  equal  vol- 
umes of  different  gases  at  the  same  tempera- 
ture and  pressure  (Avogadro's  hypothesis). 


The  molecules  themselves  are  incompressible. 


The    tendency    to    cohesion  becomes  evident 
under  some  circumstances. 


Critical  Phenomena.  —  We  may  now  use  the  terms  of  the  kin- 
etic-molecular hypothesis  in  describing  a  property  of  gases  not  yet  dis- 
cussed. When  the  concentration  of  a  gas  at  ordinary  temperatures  is 
greatly  increased  by  compression,  the  cohesive  forces  have  an  oppor- 
tunity to  produce  liquefaction.  In  many  cases,  as  with  sulphur  diox- 
ide and  carbon  dioxide,  when  the  approximation  of  the  molecules  has 
reached  a  certain  point,  the  liquid  begins  to  form  on  the  sides  of  the 
vessel.  The  condition  is  then  exactly  the  same  as  that  of  aqueous 
vapor  and  water  (p.  117),  and  no  further  increase  in  pressure  is  required 
to  complete  the  liquefaction  of  the  whole.  The  only  difference  be- 
tween steam,  at  a  pressure  below  the  aqueous  tension  of  water  at  10°, 
and  carbon  dioxide  at  the  same  temperature,  is  that  not  more  than  9.1 
mm.  of  pressure  is  required  to  liquefy  the  steam,  while  about  50  atmos- 
pheres are  needed  to  liquefy  the  carbon  dioxide. 

There  are  some  gases  in  which,  at  the  ordinary  temperature,  even 
with  the  closest  approximation  of  the  molecules,  the  cohesion  is  unable 
to  overcome  the  motion  of  the  molecules  and  draw  the  material  together 
into  the  more  compact  liquid  form.  Such  gases  are  hydrogen,  oxygen, 
nitrogen,  and  air,  which  is  a  mixture  of  the  last  two.  The  remedy  is 
obvious.  We  know  of  no  way  to  increase  the  intrinsic  cohesiveness  of 


134  INORGANIC   CHEMISTRY 

the  material,  but  we  can  reduce  the  kinetic  energy  of  the  molecules  by 
lowering  the  temperature  of  the  gas.  When  this  has  been  done  suffi- 
ciently, compression  is  followed  by  liquefaction.  Now  it  is  found  that 
there  is  a  critical  value  for  each  individual  gas  to  or  beyond  which  the 
kinetic  energy  must  be  reduced  by  lowering  the  temperature,  before 
the  cohesive  tendency  of  that  particular  gas  can  become  effective  to  pro- 
duce liquefaction.  The  highest  temperature  below  which  liquefaction 
is  possible  is  called  the  critical  temperature.  For  oxygen  this  temper- 
ature is  —118°,  for  hydrogen  about  —234°,  for  nitrogen  —146°.  For 
carbon  dioxide  it  is  31.35°,  for  sulphur  dioxide  156°,  for  water  358°. 
The  temperature  of  a  room  being  below  the  critical  point  of  the  last 
three  substances,  they  are  all  liquefiable  without  cooling,  and  more 
easily  the  farther  the  ordinary  (say  20°)  lies  below  the  critical  tem- 
perature. 

The  foregoing  is  an  example  of  how  the  images  furnished  by  an 
hypothesis  assist  us  in  understanding  facts,  and  how  the  figurative 
language  it  suggests  enables  us  to  explain  the  relations  of  the  facts 
briefly  and  vividly.  In  interpreting  statements  like  the  above,  how- 
ever, we  must  carefully  distinguish  between  the  facts  they  contain, 
which  are  permanent,  and  the  fictions  and  figures  of  speech  in  which 
we  have  clothed  them.  The  wrappings  will  doubtless  go  out  of  fashion 
and  change  in  the  future,  as  they  have  done  in  the  past,  in  response  to 
our  preference  in  emphasizing  some  different  inter-relation  in  the 
growing  body  of  truth.  The  kinetic  hypothesis  is  already  considered 
out  of  fashion  in  some  quarters. 

Kinetic  Hypothesis  Applied  to  Liquids.  —  The  phenomena  con- 
nected with  surface  tension,  such  as  coherence  into  drops,  show  that 
cohesion  plays  a  larger  part  in  liquids  than  in  gases.  The  formation 
of  vapor  from  cold  liquids,  however,  requires  us  to  suppose,  using  the 
terms  of  the  hypothesis,  that  motion  of  the  molecules  has  not  been 
annihilated  by  cohesion.  To  be  consistent,  we  have  also  to  imagine 
that  the  vapor  above  the  liquid,  for  example  the  water  in  the 
barometer  tube  in  Fig.  39  (p.  116),  is  not  composed  of  the  same 
set  of  molecules  one  minute  as  it  was  during  the  preceding  minute. 
Their  motions  must  cause  many  of  them  to  plunge  into  the  liquid, 
while  others  emerge  and  take  their  places.  When  the  water  is  first 
introduced,  there  are  no  molecules  of  vapor  in  the  space  at  all,  so  that 
emission  from  the  water  predominates.  The  pressure  of  the  vapor 
increases  as  the  concentration  of  the  molecules  of  vapor  becomes 


THE   KINETIC-MOLECULAR   HYPOTHESIS  135 

greater,  hence  the  mercury  column  falls  steadily.  At  the  same  time 
the  number  of  gaseous  molecules  plunging  into  the  water  per  second 
must  increase  in  proportion  to  the  degree  to  which  they  are  crowded  in 
the  vapor.  Hence  the  rate  at  which  vapor  molecules  enter  the  water 
must  eventually  equal  that  at  which  other  molecules  leave  the  liquid. 
At  this  point,  occasion  for  visible  change  ceases  and  the  mercury  comes 
to  rest.  We  are  bound  to  think,  however,  of  the  exchange  as  still 
going  on,  since  nothing  has  occurred  to  stop  it.  The  condition  is  not 
one  of  rest  but  of  rapid  and  equal  exchange.  Such,  described  in 
terms  of  the  hypothesis,  is  the  state  of  affairs  which  is  characteristic 
of  a  condition  of  equilibrium  (p.  117).  The  condition  is  kinetic,  and 
not  static. 

To  get  a  clear  notion  of  a  state  of  equilibrium,  we  must  distinguish 
two  opposing  tendencies,  which  when  equilibrium  is  reached  balance 
one  another.  In  the  barometer  at  rest,  for  example,  there  are  the  two 
pressures,  that  of  the  air  and  that  of  the  mercury.  Here,  one  is  the 
hail  of  molecules  leaving  the  liquid  which  is  constant  throughout  the 
above  experiment.  It  represents  the  vapor  tension  of  the  liquid 
(p.  116).  The  other  is  the  hail  of  returning  molecules  which  increases 
steadily,  at  first,  as  the  concentration  of  the  vapor  becomes  greater. 
This  is  the  vapor  pressure  of  the  vapor.  These  have  the  effect  of  op- 
posing pressures,  and  when  the  latter  becomes  equal  to  the  former, 
equilibrium  is  established.  Both  are  still  at  work,  but  neither  can 
effect  any  visible  change  in  the  system.  The  kinetic  point  of  view  we 
have  here  used  is  employed  so  continually  in  chemistry  that  it  should 
be  studied  attentively. 

When  the  temperature  of  a  liquid  is  raised,  the  kinetic  energy  of 
its  molecules  is  increased,  the  rate  at  which  they  leave  its  surface 
becomes  greater,  the  vapor  tension  increases,  and,  hence,  a  greater 
concentration  of  vapor  can  be  maintained.  This  puts  into  the  lan- 
guage of  our  hypothesis,  the  change  in  vapor  tension  with  tempera- 
ture (p.  116). 

When  the  liquid  is  placed  in  an  open  shallow  vessel,  there  is  prac- 
tically no  return  of  the  emitted  molecules.  Hence  complete  evapora- 
tion takes  place.  Elevation  of  the  temperature  hastens  the  process. 
A  draft  insures  the  total  prevention  of  all  returns,  and  has  therefore 
the  same  effect.  The  two  methods  of  assisting  the  forward  displace- 
ment of  an  equilibrium,  and  particularly  the  second,  in  which  the 
opposed  process  is  weakened  and  the  forward  process  triumphs  solely 
on  this  account,  should  be  considered  attentively  (see  Chap.  xv). 


136 


INORGANIC   CHEMISTRY 


Diffusion  in  Liquids. — We  are  further  compelled,  in  applying  our 
hypothesis,  to  suppose  that  there  is  motion  of  the  molecules  inside  the 
liquid.  When  alcohol,  as  the  lighter  liquid,  is  floated  upon  water  in  a 
cylinder  (Fig.  47),  the  plane  separating  the  liquids  is  at  first  easily 
visible.  But  soon  it  becomes  obliterated.  The  water  diffuses  upward, 
and  the  alcohol  downward,  each  sifting  its  way  through  the  other  in 
spite  of  gravity.  The  complete  mixing  of 
the  liquids  takes  a  much  longer  time  than 
in  the  case  of  two  gases.  It  may  take  months. 
But  even  here  the  hypothesis  helps  us  by 
pointing  to  the  vast  impediment  which  the 
close  packing  of  the  molecules  must  place  in 
the  way  of  the  progress  of  any  one  mole- 
cule. Further,  once  the  mixture  is  formed, 
no  tendency  to  spontaneous  separation  is 
ever  observed.  Here  again,  the  hypothesis 
shows  that  none  is  to  be  expected.  If  it  oc- 
curred, it  would  be  immediately  undone  by 
diffusion. 

Kinetic  Hypothesis  Applied  to  Solids. 

— The  properties  of  solids  differ  from  those 
of  liquids  chiefly  in  the  fact  that  the  solid 
has  a  definite  form  of  which  it  can  be  de- 
prived only  with  difficulty.  This  we  may 

explain  in  accordance  with  the  kinetic  hypothesis  by  the  supposition 
that  the  cohesion  in  solids  is  very  much  more  prominent  than  in 
liquids.  We  obtain  solids  from  liquids  by  cooling  them;  in  other  words, 
by  diminishing  the  kinetic  energy  and  therefore  the  velocity  of  the 
particles.  The  cohesive  tendency  of  the  latter  is  thus  able  to  make 
itself  felt  to  a  greater  extent.  If,  conversely,  we  heat  a  solid,  or, 
according  to  the  hypothesis,  if  we  increase  the  speed  with  which  the 
particles  move,  the  body  first  melts  and  gives  a  liquid,  and  this  finally 
boils  and  becomes  a  gas.  The  intrinsic  cohesion  of  the  particular  sub- 
stance can  undergo  no  change,  but  the  increasing  kinetic  energy  of  the 
particles  steadily  and  continuously  obliterates  its  effects.  Yet  some 
motion  still  survives  in  a  solid.  Thus  we  find  that  when  the  layer  of 
silver  is  stripped  from  a  very  old  piece  of  electroplate  the  presence  of 
this  metal  in  the  German  silver  or  copper  basis  of  the  article  is  easily 
demonstrated. 


PIG.  47. 


THE    KINETIC-MOLECULAR   HYPOTHESIS  137 

Roberts  Austen  has  found  that  if  bars  of  lead  are  prepared,  in  one  end  of  which 
an  alloy  containing  a  certain  proportion  of  gold  has  been  used,  while  the  remainder 
of  the  bar  is  composed  of  pure  lead,  the  gold  has  a  tendency  to  wander  slowly  into 
the  pure  lead.  The  process  is  greatly  aided  by  keeping  the  bars  at  a  fairly  high 
temperature,  but  one  much  below  the  melting-point  is  amply  sufficient.  After  a 
suitable  interval  of  time  the  bar  may  be  sawn  into  fragments  of  equal  length,  and 
its  parts  analyzed.  The  quantity  of  gold  in  a  section  is  found  to  increase  as  we 
approach  the  portion  of  the  bar  to  which  originally  the  whole  of  the  gold  was 
confined. 

The  tendency  of  all  solids  to  assume  crystalline  forms,  which  show 
definite  cleavage  and  other  evidences  of  structure,  distinguishes  them 
sharply  from  liquids.  The  force  of  cohesion  in  liquids  is  exercised 
equally  in  different  directions.  In  solids  it  must  differ  in  different 
directions  in  order  that  structure  may  result.  Since  each  substance 
shows  an  individual  structure  of  its  own,  these  directive  forces  must 
have  special  values  in  magnitude  and  direction  in  each  substance. 

A  crystal  arises  by  growth.  When  the  process  is  watched,  as  it 
occurs  in  a  melted  solid  or  an  evaporating  solution,  the  slow  and  sys- 
tematic addition  of  the  material  in  lines  and  layers,  as  if  according  to  a 
regular  design,  is  one  of  the  most  beautiful  and  interesting  of  natural 
phenomena.  The  fern-like  patterns  produced  by  ice  on  a  window-pane 
show  the  general  appearance  characteristic  of  crystallization  in  a  thin 
layer.  A  larger  mass  in  a  deep  vessel  gives  forms  which  are  geometri- 
cally more  perfect.  From  its  very  incipiency  the  crystal  has  the  same 
form  as  when,  later,  its  outlines  can  be  distinguished  by  the  eye.  Hence 
the  outward  form  is  only  an  expression  of  a  specific  internal  structure 
which  the  continual  reproduction  of  the  same  outward  form  on  a  larger 
and  larger  scale  leaves  as  a  memorial  of  itself  in  the  interior. 

Crystal  Forms. — Crystalline  form  is  so  continually  used  in  iden- 
tifying (p.  36)  the  substances  produced  in  chemical  actions  that  a  list  of 
the  kinds  of  forms  which  occur  will  assist  in  giving  definite  meaning 
to  our  descriptions. 

The  classification  of  crystalline  forms  is  carried  out  aqcording  to 
the  degree  of  symmetry  of  the  crystals.  Thirty-two  distinct  classes 
are  distinguished,  but  for  our  purpose  a  rougher  division  into  six 
groups  will  suffice.  These  groups  are  known  by  the  following  names: 

1.  Regular  system. 

2.  Quadratic,  or  square  prismatic  system. 

3.  Hexagonal  system. 

4.  Rhombic  system. 


138 


INORGANIC   CHEMISTRY 


FlQ.  48. 


FIG  49. 


5.  Monosymmetric,  or  monoclinic  system. 

6.  Asymmetric,  or  triclinic  system. 

The  regular  system  presents  the  most  symmetrical  figures  of  all. 
Some  forms  which  commonly  occur  are  the  octahedron 
(Fig.  48)  shown  by  alum,  the  cube  (Fig.  6,  p.  13) 
affected  by  common  salt,  and  the  dodecahedron 
(Fig.  49)  frequently  assumed  by  the  garnet. 

The  square  prismatic  system  in- 
cludes less  symmetrical  forms  than 
the  previous  one,  since  the  crystals 
are  lengthened  in  one  direction.    Fig. 
50  shows  the  condition  in  which  zir- 
con (ZrSi04),   which  furnishes  us  with  the  basis  of 
certain  incandescent  illuminating  arrangements,  occurs 
in  nature.     The  form   of   ordinary  hy- 
drated  nickel  sulphate  (NiS04,  6H20)  is  similar  to  this. 
The  hexagonal  system,  like  the  preceding,  frequently 
exhibits  elongated  prismatic  forms,  but  the  section  of 
the  crystals  is  a  hexagon  instead  of  a  square,  and  the 
termination  is  a  six-sided  pyramid.     Quartz  (Fig.  51),  or 
rock  crystal,  is  the  most  familiar  min- 
eral in  this  system.     Calcite  (CaC03), 
which    is   chemically    identical   with 
chalk,  or  marble,  takes  forms  known  as  the  scaleno- 
hedroii  (Fig.  52)  and  rhombohedron 
(Fig.  9,  p.  14),  which  are  classified  in 
a  subdivision  of  this  system.    Indeed, 
recently  it   has  become   common  to 
erect  this  into  a  separate  system  (the 
trigonal),  in  which  both  quartz  and 
calcite  are  included. 

The  rhombic  system  includes  the 
natural  forms  of  the   topaz,  and  of 
sulphur  (Fig.  1,  p.  11),  as  well  as  that 
of  potassium  permanganate  (Fig.  53),  potassium  nitrate 
(Fig.  98),  and  many  other  substances.     These  crystals 
FIG.  52.          exhibit  a  good  deal  of  symmetry,  but  their  section  is 

always  rhombic,  and  hence  the  name. 

The  monosymmetric  system  exhibits  forms  which   have  but  one 
plane  of  symmetry.     Gypsum  (Fig.  54),  which  is  hydrated  calcium 


FIG.  90. 


FIG.  51. 


THE   KINETIC-MOLECULAR   HYPOTHESIS 


139 


FIG.  53. 


sulphate  (CaS04,  2H2O),  and  felspar  are  minerals  possessing  forms  of 
this  kind.  Tartaric  acid,  rock  candy  (Fig.  55),  potassium  chlorate,  and 
hydrated  sodium  carbonate  (washing  soda)  belong  to  this  system. 

The    asymmetric  system   includes   forms  

which  have  no  plane  of  symmetry  whatever. 
Blue  vitriol  (Fig.  41,  p.  120),  (CuSO4,  5H20), 
is  one  of  the  most  familiar  substances  of  this 
kind. 

The  forms  of  crystals  which  we  may  actually  make  seldom  corre- 
spond exactly  with  the  figures.  If  we  allow  a  crystal  to  grow  upon 
the  bottom  of  a  vessel,  for  example,  it  will  usually 
have  a  tendency  to  spread  itself  out  parallel  to  the 
surface  of  the  glass,  and  when  taken  up  for  examina- 
tion will  be  found  to  present  a  somewhat  distorted 
form.  By  changing,  at  frequent  intervals,  the  face 
on  which  the  crystal  stands,  however,  uniform  growth 
in  all  directions  is  secured.  Hanging  a  small 
crystal  by  a  thread  insures  almost  ideal  develop- 
ment. Yet  the  form  even  of  distorted  crystals  can 
readily  be  recognized  by  suitable  means.  The 
shape  of  the  faces  may  indeed  be  extremely  mis- 
leading. We  find,  however,  that  the  angles  at  which 
the  faces  meet  are  always  the  same,  whatever  dis- 
proportionate growth  may  have  occurred  in  the  development  of  the 
crystal. 

Since,  in  general,  each  substance  has  a  form  of  its  own,  no  other 
substance,  as  a  rule,  can  be  used  even  partially  in  building  up  the 
crystal  (see,  however,  Isomorphism).  This 
fact  is  taken  advantage  of  in  order  to  separate 
chemical  substances  from  impurities.  The 
impure  body  is  first  dissolved  in  some  solvent. 
The  preponderating  substance  in  the  mixture, 
unless  it  is  very  much  more  soluble  than 
the  impurity,  will  then  usually  give  pure 
crystals  while  the  foreign  body  remains  in  solution. 

The  shapes  of  gems  must  not  be  confused  with  crystalline  forms. 
The  original  crystals  are  cut  and  polished  to  a  new  form  specially 
adapted  to  increase  the  ornamental  value  of  the  stone  by  causing  it  to 
reflect  more  light  (see  Diamond).  Again,  glass  is  really  a  very  viscous 
fluid  (amorphous  body,  p.  123),  and  has  no  structure  or  form  of  its  own. 


FIG.  54. 


FIG.  55. 


140  INORGANIC    CHEMISTRY 

The  word  "  crystal "  applied  to  cut  glass  would  therefore  be  misleading 
if  taken  literally. 

Crystal  Structure.  —  As  the  above  would  lead  us  to  expect,  the  study  of 
crystallized  substances  shows  that  their  peculiarities  are  not  confined  to  the  out- 
side layer.  The  outline  represents  a  certain  structure  which  permeates  the  whole 
mass.  In  crystals  of  the  regular  system,  many  of  the  ordinary  physical  properties 
are  the  same  as  those  of  an  amorphous  substance,  like  glass.  For  example,  if  we 
turn  a  sphere  out  of  crystallized  salt  and  hang  it  in  pure  water,  we  find  that  solu- 
tion takes  place  at  a  uniform  rate  all  over  the  surface.  This  is  not  the  case,  however, 
with  substances  from  any  of  the  other  systems.  Spheres  cut  from  substances  be- 
longing to  the  second  and  third  systems  would  dissolve  more,  or  less,  rapidly  in  the 
direction  of  the  chief  axis  than  in  any  other  direction,  and  so  ellipsoids  of  revolu- 
tion would  quickly  be  produced.  In  the  other  three  systems,  more  complex  forms 
would  result. 

The  tenacity  of  crystals,  to  whatever  system  they  may  belong,  is  different  in 
different  directions.  Thus,  a  crystal  of  salt  has  a  cleavage  parallel  to  any  of  the 
faces  of  the  cube,  and,  therefore,  splits  most  easily  in  one  of  three  directions  at 
right  angles  to  each  other.  Calcite,  whatever  the  outward  form  of  the  crystal, 
always  cleaves  so  as  to  give  a  rhombohedron.  Fluorite  (CaF2),  although  almost 
always  cubical  in  form,  splits  when  broken  so  as  to  give  an  octahedron. 

The  behavior  of  crystals  towards  light  is  also  extremely  interesting.  The  rate 
at  which  light  moves  through  crystals  of  the  regular  system  is  the  same  in  all 
directions.  In  other  crystals,  however,  we  find  that  it  moves  with  a  different 
speed  in  different  directions,  the  variations  in  speed  being  likewise  related  to  the 
outward  form.  Finally,  if  a  thin  slab  of  rock  salt  is  covered  with  wax  and  the 
point  of  a  heated  cone  of  metal  is  placed  in  the  center,  the  wax  melts  uniformly 
in  a  circle  around  the  point,  indicating  that  the  heat  is  conducted  with  equal  speed 
in  all  directions.  The  way  in  which  the  slab  has  been  cut  with  reference  to  the 
surface  of  the  crystal,  in  the  case  of  substances  of  the  first  of  the  above  systems, 
has  110  effect  upon  this  result.  In  all  other  cases,  however,  the  zone  of  melting 
wax  is  in  general  elliptical,  or  even  more  complex  in  form,  according  to  the  system 
to  which  the  substance  belongs  and  the  direction  in  which  the  slab  has  been  cut. 

Molecular  Magnitudes.  —  Taking  the  properties  of  matter  on 
the  one  hand  and  the  assumptions  of  the  molecular  hypothesis  on  the 
other,  we  can  estimate  the  degree  of  mutual  proximity  which  must  be 
assigned  to  the  molecules  in  order  that  the  hypothesis  may  be  consist- 
ent with  the  facts.  By  considering  the  thickness,  or  rather  the  thin- 
ness, of  thin  films,  certain  properties  of  gases,  and  other  phenomena, 
fairly  coincident,  independent  values  have  been  obtained.  According 
to  Lord  Kelvin,  1  c.c.  of  gas  at  0°  and  760  mm.  contains  not  less  than 
1020  (that  is,  1  followed  by  twenty  ciphers)  molecules.  A  globe  of 
water  as  large  as  a  football,  if  magnified  to  the  size  of  the  earth,  the 
component  molecules  being  magnified  at  the  same  time,  would  show  a 


THE   KINETIC-MOLECULAR   HYPOTHESIS  141 

structure  coarser  than  a  heap  of  small  shot  but  less  coarse  than  a 
heap  of  footballs.  The  former  estimate,  in  regard  to  gases,  leads  by 
calculation  to  the  conclusion  that  one  molecule  of  hydrogen  weighs 
0.9  x  10"24  g.,  or  9  over  1  followed  by  twenty-five  ciphers  of  a  gram. 
The  speed  of  a  hydrogen  molecule  between  collisions  must  be  about 
1840  meters  per  second.  That  of  an  oxygen  molecule  must  be  one- 
quarter  of  this,  being,  according  to  the  .law  of  diffusion  (p.  108), 
inversely  as  the  square  roots  of  the  densities. 

Formulative  and  Stochastic  Hypotheses.  —  The  nature  and 
use  of  an  hypothesis  like  the  one  we  have  been  discussing  will  now  be 
evident.  It  is  a  structure  existing  in  the  imagination.  It  cannot 
exist  anywhere  else,  because  it  includes  novelties  like  perfectly  elastic 
bodies  in  perpetual  motion.  In  making  it,  we  are  well  aware  of  the 
inverity  of  some  of  its  elements.  But  then,  as  will  have  been 
observed,  we  do  not  attempt  to  verify  the  hypothesis  itself.  We  did 
not  make  it  in  order  to  have  before  us  the  actual  structure  of  matter, 
but  in  order  to  have  a  sort  of  mechanical  moving  diagram  which 
should  assist  us  in  following  the  behavior  of  matter.  It  is  like  a  scaf- 
folding, constructed  to  enable  us  to  examine  or  work  upon  a  difficultly 
accessible  part  of  a  building,  which  we  never  for  a  moment  think  of 
as  being  a  part  of  the  building.  It  is  a  sort  of  formula.  The  algebraic 
formula  represents  magnitudes  ;  the  geometrical,  directions  and  dimen- 
sions. The  formula  in  physics,  by  the  use  of  mathematical  conven- 
tions, pictures,  for  example,  some  mode  of  behavior  of  matter.  For 
so  concrete  a  subject,  however,  the  mathematical  mode  of  expression 
is  intensely  abstract.  And  so  a  representation  in  terms  of  mechanism, 
which  is  still  a  formula,  is  frequently  resorted  to.  The  molecular 
hypothesis  is  therefore  a  formula  consisting  of  imaginary  machinery. 
Its  object  is  simply  to  help  us  in  organizing  or  formulating  knowledge 
on  a  certain  subject.  Hence  we  name  it  a  formulative  hypothesis. 

A  formula,  in  the  very  general  sense  in  which  we  have  used  the  word  above, 
is  anything  which  has  certain  properties  of  the  nature  of  form  in  common  with 
some  other  thing  or  relation  of  things.  For  example,  the  plan  of  a  city  and  a 
blackboard  diagram  are  formulae.  In  chemistry,  a  "model"  of  some  carbon 
compound,  made  up  of  balls  of  wood  and  wires,  is  just  as  much  a  formula  as  are 
the  written  symbols  showing  the  constitution  of  the  substance.  It  is  intended  to 
explain  the  behavior  of  the  substance.  From  this  sort  of  mechanical  formula  it  is 
but  a  step  to  one  consisting  of  moving  particles  and  intended  to  explain  the  be- 
havior of  gases,  The  fact  that  this  model  cannot  actually  be  constructed,  because 
perfectly  elastic  materials  are  not  available,  does  not  alter  the  ease  one  whit, 


142  INORGANIC  CHEMISTRY 

The  value  of  the  present  formulative  hypothesis,  and  of  formulative  hypothi 
in  general,  is  shown  by  the  history  of  science.  Dalton,  who  first  worked  out  a 
clear  conception  of  the  independent  behavior  of  mixed  gases  (p.  88),  used  this  hypoth- 
esis constantly  in  that  work.  Clear  understanding  of  the  separate  existence  of 
aqueous  vapor  in  the  air  could  be  reached  by  him  only  by  thinkiug  of  each  ma- 
terial as  being  made  up  of  independent  molecules.  Any  other  mode  of  conceiving 
the  mixture  that  might  readily  occur  to  one  would  involve  some  adhesion  or  inter- 
ference of  the  two  substances.  His  recognition  of  the  independent  solubilities  of 
mixed  gases,  in  proportion  to  the  partial  pressures  of  each,  would  have  been  de- 
layed or  prevented  altogether  if,  in  studying  the  results  of  his  experiments,  he  had 
not  reached  it  by  way  of  this  hypothesis. 

Other  formulative  hypotheses,  showing  the  same  contradiction  of  plain  facts 
in  their  fundamental  assumptions  that  we  have  noted  in  the  molecular  hypothesis, 
have  been  and  are  in  common  use.  The  conception  of  light  as  consisting  of  cor- 
puscles was  such  an  hypothesis  in  its  day.  The  imponderable-matter  (note  the 
contradiction)  view  of  heat  was  another.  The  undulatory  theory  of  light,  which 
postulates  a  perfectly  elastic  ether,  weightless,  frictionless,  and  lacking  every 
trace  of  impenetrability,  is  an  hypothesis  showing  the  same  inverifiable  elements. 
It  is  handled  characteristically  also,  for  we  do  not  try  by  its  means  to  learn 
more  about  ether,  but  more  about  light. 

An  essentially  different  kind  of  hypothesis  is  constantly  used  at 
every  step  in  investigation,  although  it  is  seldom  mentioned  in  books. 
When  Mitscherlich  discovered  that  Glauber's  salt  (p.  121)  gave  a  definite 
pressure  of  water  vapor,  he  at  once  formed  the  hypothesis,  that  is, 
supposition,  that  other  hydrates  would  be  found  to  do  likewise.  Ex- 
periments showed  this  supposition  to  be  correct.  The  hypothesis  was 
at  once  displaced  by  the  fact.  This  sort  of  hypothesis  predicts  the 
probable  existence  of  certain  facts  or  connections  of  facts,  hence,  reviv- 
ing a  disused  word,  we  call  it  a  stochastic  hypothesis  *  (Gk.  o-Toxao-riKo?, 
apt  to  divine  the  truth  by  conjecture).  It  differs  from  the  other  kind 
in  that  it  professes  to  be  composed  entirely  of  verifiable  facts  and 
is  subjected  to  verification  as  quickly  as  possible.  In  the  case  of 
a  formulative  hypothesis  we  have  no  expectation,  or  at  best  a  very  re- 
mote one>  of  verifying  the  hypothesis,  because  many  of  its  essential 
constituents  are  contrary  to  experience.  At  all  events,  our  efforts  are 
bent,  not  to  verifying  the  hypothesis  itself,  but  to  verifying  the  rela- 
tions between  facts  which  it  suggests. 

We  may  define  a  formulative  hypothesis  as  follows :  A  structure,  the 
essential  parts  of  which  are  assumed  facts  or  connections  of  facts  more  or  less 
inconsistent  with  known  facts,  used  in  formulating  other  known  facts.  It 
achieves  this  by  virtue  of  certain  logical  and  formal  correspondences  which  exist 

*  The  author  owes  to  Professor  Paul  Shorey,  head  of  the  department  of  Greek 
in  the  University  of  Chicago,  the  suggestion  of  this  word, 


THE   KINETIC-MOLECULAR   HYPOTHESIS  143 

between  its  abstract  qualities  and  those  of  the  facts  it  is  employed  to  explain. 
The  verification  of  the  assumed  facts  is  not  in  question,  since  their  inverity  is 
one  of  the  premises,  but  that  of  the  relations  between  the  ascertained  facts  which 
emerge,  is  the  step  to  which  the  making  of  the  hypothesis  was  only  a  preliminary. 
This  sort  of  hypothesis,  therefore,  is  a  sort  of  formula,  or  has  the  properties  of  a 
formula. 

A  stochastic  hypothesis  is  :  A  supposition  of  the  existence  of  certain  facts, 
or  connections  of  facts.  It  is  reached  by  deduction  from  other  facts  or  con- 
nections of  facts  already  ascertained,  and  is  intended  to  be  subjected  to  ultimate 
verification. 

,  The  word  theory  is  applied  to  both  kinds  of  hypotheses.  A  study  of  the 
usages  of  the  word  and  of  the  definitions  given  by  the  best  authorities  does  not 
reveal  the  existence  of  any  essential  qualitative  difference  between  a  theory  and 
an  hypothesis,  so  that  a  sharp  distinction  cannot  be  drawn  between  them.  To 
avoid  confusion,  we  shall  use  the  word  "  theory"  as  a  rule  only  in  the  sense  in  which 
it  occurs  in  phrases  like  "the  theory  of  heat,"  or  "theory  and  practice."  Here 
it  refers  to  an  aggregate  of  conceptions  and  largely  or  wholly  verified  generaliza- 
tions and  laws  which  constitute  the  abstract  (as  distinct  from  the  concrete)  state- 
ment of  the  content  of  some  branch  or  phase  of  knowledge,  so  far  as  this  has 
undergone  successful  organization. 

Lest  it  should  be  imagined  that  the  discovery  of  a  sufficiently  extensive  and 
complete  correspondence  between  a  formulative  hypothesis  and  the  facts  with 
which  it  deals  constitutes,  ipso  facto,  a  proof  of  the  verity  of  the  hypothesis  itself, 
it  must  be  stated  explicitly  that  such  an  inference  is  entirely  illogical.  Yet  this 
fallacy  is  one  of  the  commonest  into  which  the  student  of  science  falls.  This 
inference  would  be  equivalent  to  asserting  the  impossibility  of  devising  any  other 
hypothesis  which  should  correspond  equally  well  with  the  facts.  Apart  from  the 
formal  illogicality  of  such  an  inference,  experience  has  taught  us  that  sooner  or 
later  we  always  encounter  some  feature  in  the  behavior  of  the  subject  of  the 
hypothesis  which  is  different  from  that  which  the  hypothesis  would  have  led  us  to 
expect.  Thus  the  corpuscular  hypothesis  to  explain  light,  after  long  and  useful 
service,  had  finally  to  be  discarded  (see  also  Atomic  hypothesis). 

The  philosophy  of  the  molecular  hypothesis  is  discussed  by  Stallo,  Concepts 
and  Theories  of  Modern  Physics,  Chap,  vii ;  Pearson,  Grammar  of  Science,  Chap, 
vii;  James  Ward,  Naturalism  and  Agnosticism,  Lectures  iv  and  v,  and  many 
other  writers. 

The  value  of  the  kinetic-molecular  conception  as  a  formulative 
hypothesis  has  been  illustrated  in  connection  with  critical  phenomena, 
and  will  appear  again  when  we  deal  with  solutions  and  chemical  equi- 
librium. But  indeed  its  field  of  application  is  coextensive  with  the 
science  itself. 

The  conceptions  of  the  kinetic  hypothesis  are  of  especially  great  assistance  in 
rationalizing  chemical  manipulation  and  so  hastening  the  acquisition  of  an  intelli- 
gent control  of  it.  The  constructive  imagination,  on  the  use  of  which  experi- 
mental work  depends  so  much  for  its  success,  must  have  something  to  work  with, 
and  this  hypothesis  furnishes  a  tool  such  as  it  requires,  Methods  taught  by  rule 


144  INORGANIC   CHEMISTRY 

of  thumb  are  slowly  learned  and  constantly  fail  in  application.  We  may  be  told 
a  dozen  times  that  using  reagents  in  finely  powdered,  or  metals  in  granulated  con- 
dition hastens  all  interactions,  and  still  never  think  of  this  abstraction  when 
working.  But  if  it  is  suggested  that,  in  terms  of  the  kinetic  hypothesis,  molecules 
must  meet  freely  in  the  same  medium  to  react  easily,  and  that,  therefore,  the 
larger  the  surface  the  more  copious  will  be  the  supply  of  molecules  dissolving,  we 
are  likely  to  form  a  conception  of  the  reason  for  the  procedure  that  will  be 
lasting. 

When  a  student  is  told  to  concentrate  a  solution  and  set  it  aside  to  crystallize, 
why  does  he  evaporate  the  liquid  rapidly  to  dryness,  and  still  expect  the  residue 
to  appear  in  large,  well-formed  crystals  ?  Because  he  has  no  notion  of  the  necessary 
slowness  of  the  process.  But  if  he  has  the  idea  that  it  is  like  building  a  house, 
one  stone  at  a  time,  and  that  there  are  far  more  units  to  be  laid  down  according  to 
plan  in  making  the  smallest  visible  crystal  than  there  are  bricks  in  building  the 
largest  factory,  his  procedure  may  promptly  become  more  rational. 


CHAPTER   X 
SOLUTION 

WE  have  frequently  made  use  of  the  fact  that  certain  substances 
form  with  others  homogeneous  systems  which  we  call  solutions. 
Sometimes  this  property  is  taken  advantage  of  for  separating  ma- 
terials, as  in  the  case  of  the  removal  of  sulphur  from  admixture  with 
iron  and  ferrous  sulphide  (p.  11).  In  other  cases  we  carry  out  the 
interaction  of  chemical  substances,  by  first  dissolving  them  in  some 
liquid  and  then  mixing  the  solutions.  The  liquid,  commonly  water,  is 
used  as  a  vehicle  for  one  or  more  of  the  substances,  and  takes  no  part 
in  the  chemical  change.  Thus  some  knowledge  of  the  properties  of 
solutions  is  absolutely  necessary  in  order  that  we  may  employ  them 
intelligently.  In  what  follows,  we  shall  give  a  preliminary  account  of 
some  of  the  simpler  facts  about  solution. 

General  Properties  of  Solutions.  —  A  solid  may  be  distributed 
through  a  liquid  either  by  being  simply  suspended  (p.  113)  in  the 
latter  (mixture)  or  by  being  dissolved  in  it  (solution).  Similarly  a 
liquid  may  be  suspended  in  droplets  in  another,  as  in  milk  (emulsion), 
or  it  may  be  dissolved.  It  is  usually  easy  to  distinguish  between  the 
two  cases,  for  a  suspended  substance  settles  or  separates  sooner  or 
later,  while  a  dissolved  one  shows  no  such  tendency.  The  cases  are 
exceptional  where  the  subdivision  of  a  suspended  substance  is  so 
minute  (colloidal  solution)  as  to  make  its  retention  by  filter  paper 
impossible.  If  a  liquid  is  opalescent  or  opaque,  then  we  have  a  case 
of  suspension.  A  solution  is  a  clear,  transparent,  perfectly  homo- 
geneous liquid,  in  which  the  dissolved  substance  seems  to  have  been 
dispersed  so  completely  that  the  liquid  cannot  be  distinguished  by  the 
eye  from  a  pure  substance. 

There  is  no  limit  to  the  amount  of  dissipation  which  may  thus  be 
produced.  A  single  fragment  of  potassium  permanganate,  for  ex- 
ample, which  gives  a  very  deep  purple  solution  in  water,  may  be  dis- 
solved in  a  liter  or  even  in  twenty  liters  of  water,  and  the  purple 
tinge  which  it  gives  to  the  liquid  will  still  be  perfectly  perceptible  in 

145 


146  INORGANIC   CHEMISTRY 

every  part  of  the  larger  volume.  The  characteristics,  therefore,  of 
solution  are  absence  of  settling,  homogeneity,  and  extremely  minute 
subdivision  of  the  dissolved  substance. 

The  Scope  of  the  Word.  —  The  word  is  used  for  other  systems 
than  those  containing  a  solid  body  dissolved  in  a  liquid.  Thus,  liquids 
also  may  be  dissolved  in  liquids,  as  alcohol  in  water.  Again,  if  we 
warm  ordinary  water,  bubbles  of  gas  appear  on  the  sides  of  the  vessel 
before  the  water  has  approached  the  boiling-point.  They  are  found 
to  be  air.  Further  study  of  the  subject  shows  that  agitation  of  any 
gas  with  water  results  in  the  solution  of  a  large  or  small  quantity  of  the 
gas,  and  heat  will  usually  drive  the  gas  out  again.  It  appears  there- 
fore that  solids,  liquids,  and  gases  can  equally  form  solutions  in  liquids. 

The  absorption  of  hydrogen  by  palladium  (at  all  events  after  a 
certain  point)  and  by  iron  takes  place  in  accordance  with  the  same 
laws  as  the  solution  of  solids  in  liquids,  and  the  results  may  be 
described  therefore  as  true  solutions.  Liquids  are  in  some  cases 
absorbed  by  solids,  and  homogeneous  mixtures  of  solids  with  solids 
are  perfectly  familiar.  The  sapphire  is  a  solution  of  a  small  amount 
of  a  strongly  colored  substance,  in  a  large  amount  of  colorless  aluminium 
oxide.  It  may  therefore  be  stated  that  solution  of  gases,  liquids,  and 
solids  in  solids  appears  to  be  possible. 

Limits  of  Solubility.  —  The  next  question  which  naturally  occurs 
to  us  is  as  to  whether  the  mingling  of  two  substances  in  this  manner 
has  any  limits.  We  find  that  the  results  of  experiment  in  this 
direction  may  be  divided  into  two  classes.  Some  pairs,  of  liquids 
particularly,  may  be  mixed  in  any  proportions  whatever.  Alcohol  and 
water  is  such  a  pair.  On  the  other  hand,  at  the  ordinary  laboratory 
temperature,  we  can  scarcely  take  a  fragment  of  marble  (CaC08)  so 
small  that  it  will  dissolve  completely  in  100  c.c.  of  pure  water.  Under 
the  same  conditions  any  amount  of  potassium  chlorate  up  to  5  g.  will 
almost  completely  disappear  after  vigorous  stirring,  while  90  g.  of 
ordinary  Epsom  salts  (hydrated  magnesium  sulphate),  but  not  more, 
may  be  dissolved  in  about  the  same  amount  of  water.  In  fact,  most 
solids  may  be  dissolved  in  a  liquid  only  up  to  a  certain  limit,  which 
with  different  solids  may  range  from  a  scarcely  perceptible  to  a  very 
large  amount.  No  substance  is  absolutely  insoluble.  But  for  the 
sake  of  brevity  we  call  marble,  for  example,  "  insoluble "  because  in 
most  connections  it  may  be  so  considered. 


SOLUTION 


147 


Recognition  and  Measurement  of  Solubility.  —  The  only 
method  of  recognizing  with  certainty  whether  a  solid  is  soluble  in  a 
liquid  or  not  is  to  filter  the  mixture  and  evap- 
orate a  few  drops  of  the  filtrate  on  a  clean 
watch-glass.  For  learning  how  much  of  the 
body  is  contained  in  a  given  solution,  a  weighed 
quantity  of  the  solution  is  evaporated  to  dry- 
ness  and  the  weight  of  the  residue  deter- 
mined. When  the  dissolved  substance  is 
volatile,  its  presence  is  often  shown  by  some 
chemical  test  (p.  99). 

Ether  and  water  is  a  case  typical  of  the 
behavior  of  two  liquids,  each  somewhat  soluble 
in  the  other.  After  being  shaken  together, 
they  seem  to  separate  again  completely  into 
two  layers  (Fig.  56)  with  the  ether  uppermost. 
If,  however,  the  water  is  withdrawn  from 
beneath  the  ether,  we  find  that,  when  heated, 
it  gives  off  quantities  of  ether  vapor  which  can 
be  set  on  fire.  Conversely,  the  addition  of 
anhydrous  cupric  sulphate  to  a  sample  of  the  FlG-  56- 

ether  shows  the  presence  of  water  in  the  latter, 

for  the  blue  hydrated  form  of  the  substance  is  at  once  produced.     In 
some  common  cases  the  maximum  solubilities  at  22°  are  as  follows  : 


SUBSTANCE. 

GRAMS  OF  SUBSTANCE 

IN   100 

GRAMS  OF  WATER. 

GRAMS  OF  WATER 
IN  100 
GRAMS  SUBSTANCE. 

Alcohol      

No  limit 

No  limit 

Ether    

2.16 

11  02 

0.64 

0  10 

1.24 

0  13 

It  must  be  stated  explicitly  that  in  going  into  solution,  as  we  have 
used  the  term,  a  compound  dissolves  as  a  whole,  and,  if  the  compound 
is  pure  (p.  34),  any  residue  has  the  same  chemical  composition  as  the 
part  which  has  dissolved. 

Terminology.  —  In  order  to  describe  the  relations  of  the  compo- 
nents of  a  solution,  certain  conceptions  and  corresponding  technical 
expressions  are  required.  It  is  customary  to  speak  of  the  substance 


148  INORGANIC   CHEMISTRY 

which,  like  water  in  most  cases,  forms  the  bulk  of  the  solution,  as  the 
solvent.  To  express  the  substance  which  is  dissolved,  the  word  solute 
is  frequently  used,  and  will  be  employed  when  we  wish  to  avoid  cir- 
cumlocution. The  term  "strength"  is  too  indefinite  for  scientific 
purposes.  It  may  imply  activity,  or  power  of  resistance,  or  pungency 
(in  an  odor),  or,  as  in  the  case  of  solutions,  it  may  be  a  measure  of 
quantity.  The  amount  of  the  substance  which  has  been  dissolved 
by  a  given  quantity  of  the  solvent  is  therefore  described  as  the 
concentration  of  the  solution.  A  solution  containing  a  small  propor- 
tion of  the  dissolved  body  is  called  dilute ;  it  has  a  small  concentration. 
One  which  contains  a  larger  amount  is  more  concentrated.  Very 
"  strong "  solutions  are  frequently  spoken  of  simply  as  concentrated 
solutions.  The  partial  removal  of  the  solvent  by  evaporation  is  called 
concentrating,  its  total  removal  evaporating  to  dryness.  Finally,  since 
there  is  a  limit  to  the  solubility  of  most  substances,  a  solution  is 
described  as  saturated  when  the  solute  has  given  as  much  material  to 
the  solvent  as  it  can.  This  state  is  reached  after  prolonged  agitation 
with  an  excess  of  the  gas,  the  liquid,  or  the  finely  powdered  solid,  as  the 
case  may  be.  The  larger  the  excess,  the  sooner  saturation  is  attained. 
The  maximum  concentration  attainable  in  this  way  is  called  the 
solubility  of  the  substance  in  a  given  solvent. 

The  distinction  between  solute  and  solvent  is  made  merely  for  convenience. 
Theoretically  there  is  no  distinction  between  the  components  of  a  solution. 

The  concentrations  of  solutions,  saturated  and  otherwise,  are  some- 
times expressed  in  physical,  and  sometimes  in  chemical,  units  of 
weight.  When  physical  units  are  employed,  as  in  the  above  table,  we 
give  the  number  of  grams  of  the  solute  held  in  solution  by  one  hundred 
of  the  solvent,  or,  occasionally,  the  number  of  grams  in  one  hundred 
of  the  solution. 

When  chemical  units  of  weight  are  employed,  two  different  plans  are 
possible,  and  both  are  in  use.  Either  the  equivalent  (p.  49)  or  the  atomic 
weights  may  be  taken  as  a  basis  of  measurement.  In  the  former  case, 
the  solutions  are  called  normal  solutions,  and  in  the  latter,  for  a  reason 
which  will  appear  later  (Chap,  xii),  molar  solutions. 

A  normal  solution  contains  one  gram-equivalent  of  the  solute  in  one 
liter  of  solution.  The  word  "equivalent"  has  been  used  hitherto  only  of 
elements,  and  this  application  of  the  expression  involves  an  extension 
of  its  meaning.  An  equivalent  weight  of  a  compound  is  that  amount 
of  it  which  will  interact  with  one  equivalent  of  an  element.  Thus,  a 


SOLUTION 


149 


formula-weight  of  hydrochloric  acid  HC1  (36.5  g.)  is  also  an  equivalent 
weight,  for  it  contains  1  g.  of  hydrogen,  and  this  amount  of  hydrogen  is 
displaceable  by  one  equivalent  weight  of  a  metal.  A  formula-weight 
of  sulphuric  acid  H2SO4  (98  g.),  however,  contains  two  equivalents  of 
the  compound,  and  a  formula-weight  of  aluminium  chloride  AlCL, 
(133.5  g.)  three  equivalents.  Hence  normal  solutions  of  these  three 
substances  contain  respectively  36.5  g.  (HC1),  49  g.  (H2S04),  and  44.5 
g.  (AlClg)  per  liter  of  solution.  The  special  property  of  normal  solu- 
tions is,  obviously,  that  equal  volumes  of  two  of  them  contain  the  exact 
proportions  of  the  solutes  which  are  required  for  complete  interaction. 
Solutions  of  this  kind  are  much  used  in  quantitative  analysis. 

Solutions  of  different  concentrations  all  prepared  on  the  above 
basis  are  named  as  follows,  and  are  often  indicated  by  the  abbreviations 
appended : 


QUANTITY  OF  SOLUTE 
PEB  LITER. 

NAME. 

ABBREVIATION. 

One  hundredth  of  one  gram-) 
eou.iva.lent     i 

Centi-normal 

iM  or  '01  N 

One  tenth  of  one  gram-equiv-  ) 
alent                             .     .       i 

Deci-normal 

N 
in   or    -1  N 

One  half  of  one  gram-equiva-  ) 
lent  1 

Semi-normal 

1U 

N 
-s-   or    AN 

One  gram-equivalent      .     .     . 
Two  and  a  half  gram-equivalents 

Normal 
Two  and  a  half  normal 

N 
2$  N  or  2.6  JV 

A  molar  solution  contains  one  mole  (gram-molecular  weight)  of  the 
solute  in  one  liter  of  solution.  When  molecular  formulae  (see  Chap,  xii)  are 
used,  this  means  one  gram-formula  weight  per  liter.  In  the  cases  cited  above,  the 
molar  solution  contains  36.5  g.  (HC1),  98  g.  (H2SO4),  and  133.6  g.  (A1C13)  per  liter. 
As  will  be  seen,  the  concentration  of  molar  and  normal  solutions  are  necessarily 
identical  when  the  radicals  are  univalent.  Other  concentrations  are  described  as 
deci-molar  (M/IQ  or  .1 M),  two  and  a  half  molar  (2.6 .M),  and  so  forth,  on  the  same 
plan  as  before. 

There  is  also  a  chemical  unit  of  volume  (see  Chap,  xii)  which  is  the  volume 
occupied  by  a  mole  (gram-molecular  weight)  of  a  gas  (or  dissolved  substance)  at 
0°  and  760  mm.  pressure  (gaseous  or  osmotic).  This  volume  averages  22.4  liters,  and 
is  called  the  gram-molecular  or  molar  volume  (G.  M.V.).  The  unit  of  concentra- 
tion for  many  theoretical  purposes  is,  therefore,  that  of  one  mole  in  22.4  liters. 

Solution  One  of  the  Physical  States  of  Aggregation  of  Matter. 

—  When  a  solid  body  dissolves  in  a  liquid,  the  properties  of  the  body 
undergo   a  very   marked  change,  which  to  all  appearance  might  be 


150  INORGANIC   CHEMISTRY 

chemical.  Yet,  besides  the  ease  with  which  a  liquid  may  be  removed 
by  evaporation  and  the  solid  recovered  unchanged,  we  note  particularly 
that  the  concentration  of  a  saturated  solution  cannot  be  expressed  in 
terms  of  integral  multiples  of  the  chemical  combining  weights.  If, 
therefore,  the  process  were  to  be  regarded  as  chemical,  several  impor- 
tant generalizations  would  have  to  be  revised  or  discarded  (cf.  p.  48). 
We  shall  see  also,  in  a  later  paragraph,  that  the  quantity  of  a  solid 
which  a  liquid  may  take  up  varies  with  the  slightest  change  in  tem- 
perature. Now  we  do  not  find  the  composition  of  chemical  compounds 
so  to  vary.  The  solution  of  a  solid  may  therefore  be  likened  to  a 
change  in  state,  such  as  the  conversion  of  a  liquid  into  a  gas  or  a  solid. 
As  there  is  danger  of  confusion  arising,  we  may  repeat  that  a  com- 
pound is  homogeneous  and  its  composition  is  expressible  in  chemical 
units  of  weight ;  a  saturated  solution  is  homogeneous  but  its  concentra- 
tion varies  with  temperature  so  that  chemical  units  cannot  be  used  to 
describe  its  composition ;  a  mixture,  as  of  two  solids  or  two  liquids,  is 
neither  homogeneous  nor  in  any  way  definite  in  composition. 

Kinetic-Molecular  Hypothesis  Applied  to  the  State  of  Solu- 
tion. —  Accepting  solution  as  a  physical  state,  we  may  now  apply  the 
same  formulative  hypothesis  to  the  explanation  of  the  behavior  of  a 
substance  in  solution  as  to  matter  in  the  gaseous  or  liquid  states.  We 
saw  that  a  solid  body,  which  is  ordinarily  condensed  in  a  small  space, 
can  be  disseminated  by  the  use  of  a  solvent  through  a  very  large  one. 
The  molecules  of  the  solid  become  scattered  like  those  of  a  gas  or  vapor 
through  a  much  greater  volume.  We  may  regard  the  dissolved  sub- 
stance as  being,  practically,  in  a  gaseous  or  g'wem-gaseous  condition. 
The  molecules  are  torn  apart  from  one  another,  their  cohesion  is  over- 
come, and  their  freedom  of  motion  is  in  a  measure  restored.  It  is  true 
that  they  could  not  continue  to  occupy  this  large  volume  for  a  moment 
in  the  absence  of  the  solvent.  But  we  may  bring  this  into  relation 
with  the  case  of  a  vapor  by  saying  that  a  solid  body,  like  common  salt, 
can  only  evaporate  appreciably  at  the  ordinary  temperature,  and 
occupy  a  large  space,  when  that  space  is  already  filled  with  a  suitable 
liquid.  The  latter  acts  as  a  vehicle  for  the  particles  of  the  solid.  A 
volatile  liquid,  on  the  contrary,  can  dissolve  in  an  empty  space  and  fill 
it  with  its  particles  without  any  vehicle  being  required. 

This  conception  of  the  quasi-gaseous  condition  of  a  dissolved  sub- 
stance would  be  simply  fantastic  if  it  did  not  lead  us  to  a  better  under- 
standing of  the  behavior  of  solutions.  Now  there  are  certain  proper- 


SOLUTION 


151 


ties  of  gases  already  discussed  which  we  should  naturally  proceed  to 
look  for  in  the  case  of  dissolved  bodies  if  this  theory  were  correct. 
These  are  the  properties  described  under  the  law  of  diffusion,  Boyle's 
law,  and  Charles'  law. 

It  is  easy  to  show  that,  if  we  place  a  quantity  of  the  pure  solvent 
(Fig.  57)  above  a  concentrated  solution  of  a  substance,  and  then  set 
the  arrangement  aside,  the 
dissolved  body  slowly  makes 
its  way  through  the  liquid 
(Fig.  58),  obliterating  the 
original  plane  of  separation. 
Eventually  the  dissolved 
body  scatters  itself  uniform- 
ly through  the  whole.  In 
other  words,  the  particles  of 
the  dissolved  substance  ex- 
hibit the  property  of  diffu- 
sion in  the  same  way  as  do 
those  of  gases. 

When  the  diffusion  of  a 
gas  is  resisted  by  a  suitable 
partition,  we  find  pressure  is 
exercised  upon  the  walls  of  FIG.  57.  FIG.  58. 

the   vessel    and    upon    the 

partition.  It  is  possible  to  show  that  the  particles  of  a  dissolved  sub- 
stance exercise  a  pressure  of  a  very  similar  kind.  This  pressure  is 
spoken  of  as  osmotic  pressure  (q.v.).  In  order  to  get  evidence  of  this 
pressure,  we  must  place  a  partition  between  the  solution  and  the  pure 
solvent  (Fig.  57),  and  we  must  choose  the  material  of  this  partition  so 
that  its  substance  gives  free  passage  to  the  solvent  while  it  resists  the 
exit  of  the  dissolved  material.  The  particles  of  the  latter,  then,  by 
their  impacts  upon  it,  produce  the  effect  of  pressure  exactly  as  in  the 
case  of  gases.  This  pressure  is  found  to  be  proportional  to  the  con- 
centration of  the  solution.  In  other  words,  it  depends  upon  the  degree 
of  crowding  of  the  molecules  of  the  dissolved  substance.  It  corre- 
sponds therefore  exactly  to  the  pressure  of  gases  in  this  respect,  and 
Boyle's  law  (p.  81)  expresses  the  relation  equally  for  both. 

It  may  be  added,  also,  that  osmotic  pressure  increases  when  a 
solution  is  warmed,  and  it  gains  ?}F  of  the  value  it  had  at  0°  for  every 
degree  through  which  the  temperature  is  raised.  In  other  words, 


152 


INORGANIC   CHEMISTRY 


Charles'  law  (p.  87)  describes  the  change  in  osmotic  pressure  with 
change  in  temperature  just  as  correctly  as  it  does  the  change  in  gaseous 
pressure. 

Kinetic-Molecular  Hypothesis  Applied  to  the  Process  of 
Solution.  —  We  may  now  apply  the  same  hypothesis  to  the  process  of 
dissolving,  with  a  view  more  especially  to  explaining  why  the  process 
of  dissolving  ceases,  in  spite  of  the  presence  of  excess  of  the  solute, 
when  a  certain  concentration  has  been  reached.  If  some  of  the  mate- 
rial dissolves,  why  not  more  ? 

Let  us  suppose  that  it  is  the  dissolving  of  common  salt  in  water 
(Fig.  59)  which  we  wish  to  explain  in  detail.  We  believe  that  in  the 
solid  substance  the  molecules  are  somewhat 
closely  packed  together,  while  in  the  solution 
they  are  rather  sparsely  distributed.  If  there 
were  no  water  over  the  salt,  practically  none 
of  the  particles  of  the  latter  would  be  able  to 
leave  the  solid  and  enter  the  space  above.  Thus, 
the  process  of  solution  must  consist  in  the 
loosening  of  the  molecules  on  the  surface  and 
their  passage  into  the  liquid.  By  diffusion,  the 
free  molecules  will  gradually  move  away  from 
the  neighborhood  of  the  surface  of  the  solid  and 
make  room  for  others,  and  thus,  if  the  system 
remains  undisturbed,  the  liquid  will  eventually 
-  become  a  solution  of  uniform  concentration.  If 
a  large  enough  amount  of  the  solid  has  been 
provided,  the  ultimate  condition  will  be  that  of 
a  saturated  solution  with  excess  of  the  solid 
beneath.  If  we  had  proper  means  of  measuring 

it,  the  tendency  of  the  molecules  to  leave  the  solid  in  the  presence 
of  a  given  liquid  would  give  the  effect  of  a  kind  of  pressure.  This  is 
spoken  of  as  solution  pressure. 

Now  the  molecules,  after  having  entered  the  liquid,  move  in 
every  direction,  and  consequently  some  of  them  will  return  to  the 
solid  and  attach  themselves  to  it.  The  frequency  with  which  this 
will  occur  will  be  greater  as  the  crowding  of  particles  in  the  liquid 
increases,  so  that  a  stage  will  eventually  be  reached  at  which  the 
number  of  molecules  leaving  the  solid  will  be  no  greater  than  that 
landing  upon  it  in  a  given  time.  If  the  whole  of  the  liquid  has  mean- 


FlG.  59. 


SOLUTION  153 

while  become  equally  charged  with  dissolved  molecules,  there  will  be  no 
chance  that  the  field  of  liquid  immediately  round  the  solid  will  lose 
them  by  diffusion,  so  that  a  condition  of  balance  or  equilibrium  will 
have  been  established  :  KaCl  (solid)  <=±NaCl  (diss'd).  The  motion  of 
the  particles  in  the  liquid  produces  what  we  have  called  osmotic  press- 
ure ;  and  when  the  osmotic  pressure,  by  the  continual  increase  in  the 
number  of  dissolved  molecules,  becomes  equal  to  the  solution  press- 
ure, increase  in  concentration  of  the  solution  ceases.  It  is  at  this 
point  that  we  speak  of  the  solution  as  being  saturated  with  respect  to 
the  particular  substance  dissolving.  The  analogy  to  vapor  tension 
and  vapor  pressure  (p.  135)  is  evident. 

The  necessity  of  distinguishing  between  the  fictions  used  in  thinking  about  and 
describing  the  phenomena  of  solution  and  the  facts  themselves  must  be  empha- 
sized here  as  it  was  in  the  preceding  chapter.  The  arrangement  and  expression  of 
the  facts  about  solutions  in  terms  of  the  kinetic-molecular  hypothesis  is  known  as 
the  theory  of  solutions. 

Independent  Solubility.  —  Just  as  two  gases,  when  mixed,  are 
independent  of  one  another  (p.  88),  and  have  severally  the  same 
pressure,  solubility,  and  so  forth,  as  they  would  possess  if  each  alone 
occupied  the  same  space,  so  is  it  with  dissolved  substances.  In  gen- 
eral, a  volume  of  water,  in  which  a  moderate  amount  of  some  substance 
has  been  dissolved,  will  take  up  as  much  of  a  second  substance  as 
would  an  equal  volume  of  pure  water.  Thus,  water  containing  some 
sugar  will  dissolve  as  much  sodium  chloride  as  the  same  amount  of 
pure  water.  In  the  point  of  view  of  the  kinetic-molecular  hypothesis, 
the  dissolved  molecules  of  sugar  have  no  connection  with,  or  influence 
upon,  the  mechanism  which  determines  the  solubility  of  the  salt, 
namely,  the  exchange  of  salt  molecules  between  the  suspended,  dis- 
solving crystals  and  the  solution. 

Naturally  this  principle  of  independent  solubility  does  not  hold 
with  any  degree  of  exactness  when  the  concentration  of  the  substance 
already  present  is  great.  It  fails  also  when  the  two  solutes  interact 
chemically,  as  will  usually  be  the  case,  for  example,  when  each  is  an 
acid,  base,  or  salt  (see  Chap.  xvi).  The  solubility  is  affected  in  an 
especial  degree  when  the  two  substances  have  one  radical  in  common, 
as  when  they  are  nitric  acid  and  a  nitrate,  or  two  chlorides  (see  Ionic 
equilibrium). 

Solution  of  a  Gas  in  a  Liquid.  —  The  same  conceptions  may  be 
used  to  explain  any  case  of  solution.  Let  us  take  that  of  oxygen  con- 


154 


INORGANIC   CHEMISTRY 


ducted  into  a  bottle  which  is  partially  filled  with  water  (Fig.  60),  no 
other  gas  being  present  in  the  space  above  the  liquid.  As  the  mole- 
cules of  the  gas  impinge  upon  the  liquid,  some  of  them  pass  into  it 
and  dissolve.  The  particles  which  have  thus  gained  access  to  the 

liquid  move  about  in  every  direc- 
tion, and,  as  they  become  more 
and  more  numerous,  a  larger  and 
larger  number  will  escape  from 
the  surface  and  pass  back  into  the 
gaseous  condition.  At  first,  this 
reaction  will  be  slight,  but  event- 
ually, as  the  solution  increases  in 
concentration,  it  must  become 
equal  in  rate  to  the  process  of 
solution  itself.  It  is  assumed 
that  the  supply  of  gas  is  main- 
tained at  a  uniform  pressure,  and 
therefore  uniform  molecular  con- 
centration, during  the  whole  pro- 
cess. Once  more  we  shall  have 
a  state  of  balance  or  equilibrium, 
and  the  liquid  will  be  saturated, 
this  time  with  a  gas  :  0  (gas)  +±  O 

(diss'd).  It  is  found,  as  the  hypothesis  would  lead  us  to  expect,  that 
the  concentration  of  the  saturated  solution  of  a  gas  is  proportional  to 
the  pressure  at  which  the  gas  is  supplied.  This  is  usually  known  as 
Henry's  law. 

Henry's  law  is  anticipated  from  the  theory  when  we  consider  that  the  pressure 
of  the  gas  is  proportional  to  the  degree  of  crowding  of  its  molecules.  On  the 
other  hand,  the  concentration  that  can  be  maintained  in  the  saturated  solution  is 
proportional  to  the  frequency  of  the  impacts  of  the  molecules  of  the  gas  upon  the 
surface  of  the  liquid,  and  this  in  turn  will  be  proportional  to  the  degree  to  which 
they  are  crowded  in  the  space  above  "the  liquid.  Hence  the  concentration  of  the 
saturated  solution  is  proportional  to  the  gaseous  pressure. 

The  solubility  of  different  gases  varies  much.  One  volume  of 
water  will  dissolve  1050  volumes  of  ammonia  at  0°  and  760  mm., 
while  it  will  dissolve  only  about  0.02  volumes  of  hydrogen  under  the 
same  conditions.  In  one  volume  of  alcohol,  at  0°  and  760  mm.,  17.9 
volumes  of  hydrogen  sulphide  or  0.07  volumes  of  hydrogen  may  be 
dissolved.  The  law  describes  the  behavior  of  the  bodies  with  exact- 


FIG.  60. 


SOLUTION  155 

ness  only  when  low  gaseous  pressures  and  gases  whose  solubility  is 
small  are  in  question.  Great  solubility  must  be  due  in  part,  probably, 
to  chemical  union  between  the  material  of  the  gas  and  the  solvent,  or 
to  cohesive  influences  which  the  molecules  of  the  dissolved  gas  exert 
upon  each  other.  The  hypothesis,  on  the  other  hand,  considers  only 
an  ideal  behavior  involving  complete  chemical  and  physical  independ- 
ence of  the  molecules. 

When  more  than  one  gas  is  in  contact  with  the  solvent,  the  hy- 
pothesis enables  us  successfully  to  foretell  what  will  happen.  The 
quantity  of  each  gas  which  can  remain  dissolved  must  depend  simply 
upon  the  frequency  with  which  its  own  molecules  strike  the  liquid, 
and  must  be  independent  of  the  presence  of  the  other  gas.  Hence  the 
solubility  of  each  gas  is  the  same  as  if  it  were  present  alone  at  its 
own  partial  pressure  (p.  88).  Dalton  used  these  very  considerations 
in  drawing  this  conclusion  from  his  experimental  data. 

Air  dissolving  in  water  is  an  illustration  of  this  principle.  It  does 
not  dissolve  as  a  whole,  but  the  oxygen  and  nitrogen  dissolve  each  in 
proportion  to  its  intrinsic  solubility  and  partial  pressure.  i/ 

It  is  easy,  by  the  use  of  this  law,  to  form  an  approximate  estimate  of  the  pro- 
portion of  oxygen  to  nitrogen  in  the  dissolved  gases.  The  air  may  be  taken  to  be 
at  760  mm.,  and  its  composition  by  volume  roughly  1/5  oxygen  and  4/5  nitrogen. 
The  separate  solubilities  of  the  gases  at  760  mm.  are,  respectively,  4  and  2 
volumes  in  100  volumes  of  water.  Their  partial  pressures  being  1/5  and  4/5  of  an 
atmosphere,  the  amounts  actually  dissolved  will  be  4  x  1/5  =  0.8  and  2  x  4/5  = 
1.6  in  100  volumes  of  water.  The  ratio  of  free  oxygen  to  nitrogen  in  the  water 
will  therefore  be  1 :  2. 

Two  Immiscible  Solvents :  Law  of  Partition.  —  An  interesting 
application  of  the  same  ideas  may  be  made  to  a  case  which  occurs 
very  commonly  in  chemical  work.  If  we  shake  up  a  small  particle  of 
iodine  with  water,  we  rind  that  it  dissolves  slowly,  giving  eventually  a 
saturated  but  very  dilute  solution.  If  now  ether  in  sufficient  quantity 
be  shaken  with  the  aqueous  solution,  the  greater  part  of  the  iodine  will 
find  its  way  into  the  ether,  and  be  contained  in  the  brown  layer  which 
rises  to  the  top.  The  process  of  removing  a  substance  partially  from 
solution  in  one  solvent  and  securing  it  in  another  is  called  extraction. 
We  find  in  such  cases  that  neither  solvent  can  entirely  deprive  the 
other  of  the  whole  of  the  dissolved  substance,  if  the  latter  is  soluble 
in  both  independently  :  I  (in  Aq.)  ^±  I  (in  ether).  The  partition  of 
the  substance  takes  place  in  proportion  to  its  solubility  in  each  sol- 
vent. It  is  found  that  any  amount  of  the  solute,  up  to  the  rnaxi- 


156  INORGANIC   CHEMISTRY 

mum  the  system  can  contain,  provided  this  does  not  involve  too  high 
a  concentration  in  either  solvent,  is  divided  so  that  the  ratio  of  the 
concentrations  in  the  two  solvents  is  always  the  same.  In  the  case  of 
iodine  divided  between  water  and  ether,  this  ratio  is  about  1 : 200. 

The  aqueous  solution  of  potassium  iodide  has  a  very  great  power  of  dissolving 
iodine,  and  we  find  that  in  the  presence  of  this  salt  the  ether  leaves  a  much  larger 
share  of  the  element  in  the  lower  layer.  A  part  of  the  iodine  combines  to  form 
KI3,  however,  so  that  this  is  not  a  case  of  simple  solution,  and  the  law  of  partition 
does  not  hold.  The  chemical  equilibrium  (see  Chap,  xv)  between  free  and  com- 
bined iodine  in  the  aqueous  layer  has  to  be  considered. 

Influence  of  Temperature  on  Solubility.  —  The  quantity  of  a 
substance  which  we  can  dissolve  in  a  fixed  amount  of  a  given  solvent 
depends  very  largely  upon  the  temperature  of  both.  Usually  the  solu- 
bility increases  with  rise  in  temperature.  Measurements  may  be  made 
by  the  method  described  before  (p.  147),  using  excess  of  the  finely  pow- 
dered solute  with  different  portions  of  the  same  solvent  in  vessels  kept 
at  different  temperatures.  The  most  useful  way  of  representing  the  re- 
sults is  to  plot  them  graphically.  The  diagram  (Fig.  61)  shows  the 
curves  for  a  few  familiar  substances.  The  ordinates  represent  the  num- 
ber of  grams  of  the  anhydrous  compound  which  is  held  in  solution  by 
100  g.  of  water  in  each  case.  The  abscissae  represent  the  tempera- 
tures. The  concentration  for  any  temperature  can  be  read  off  at  once. 
Thus  100  g.  of  water  holds  13  g.  of  potassium  nitrate  in  solution  at  0° 
and  200  g.  at  87°.  The  increase  in  solubility  is  here  enormous.  On 
the  other  hand,  the  same  quantity  of  water  will  hold  35.6  g.  of  sodium 
chloride  in  solution  at  0°  and  40  g.  at  100°.  The  difference  is  shown 
at  once  when  we  examine  the  curves  and  observe  that  the  line  repre- 
senting the  solubility  of  sodium  chloride  scarcely  rises  at  all  between 
0°  and  100°,  while  that  of  potassium  nitrate  is  extremely  steep. 

Cases  in  which  the  solubility  decreases  with  rise  in  temperature  are 
less  common.  When  cold  water  is  saturated  with  calcium  citrate,  and 
the  solution  is  then  warmed,  a  large  part  of  the  salt  is  quickly  pre- 
cipitated. 

When  triethylamine,  an  organic  base,  N(C2H5)g,  liquid  at  ordinary  tempera- 
tures, is  added  to  cold  water  until  no  more  will  dissolve,  the  solution,  which  is 
perfectly  clear  and  transparent,  on  being  warmed  with  the  hand  at  once  becomes 
clouded  from  the  separation  of  the  two  liquids.  A  comparatively  slight  elevation 
in  temperature  causes  a  separation  into  two  distinct  layers. 

Phases.  —  We  can  frequently  abbreviate  our  statements  by  using 
two  words  of  broad  significance,  one  of  which  hag  already  been  em- 


SOLUTION 


167 


50°  00°  7<>  80°  00°  100 


158 


INORGANIC   CHEMISTRY 


ployed.  A  set  of  materials  in  or  tending  towards  a  condition  of  equi- 
librium is  called  a  system.  The  discrete  parts  of  an  inhomogeneous 
system  are  called  its  phases.  Thus,  a  liquid  with  its  vapor  forms  a 
sj^stem  with  a  liquid  phase  and  a  vapor  phase.  A  saturated  solution 
is  a  system  with  three  phases,  the  undissolved  excess  of  the  solute 
(solid  phase),  the  solution,  and  the  vapor. 

Equilibrium  in  a  Saturated  Solution.  —  Once  a  solution  has  be- 
come saturated,  the  dissolving  substance  remains  thereafter  unchanged 
in  amount  no  matter  how  long  the  materials  are  left  in  contact.  In 
technical  terms,  the  quantity  of  each  phase  has  no  influence  on  the 

concentration  of  any 
of  them.  A  greater 
excess  of  the  solute 
forces  no  more  mat- 
ter into  solution  than 
does  a  small  excess. 
It  should  be 
clearly  understood 
that  the  kinetic  hy- 
pothesis requires  us 
to  assume  that  an 
exchange  of  mole^ 
cules  (p.  135)  is  still 
going  on  between  the 
solid  and  the  solu- 
tion.  That  this  con- 
ception is  correct 

may  be  shown  in  various  ways.  Thus,  if  a  crystal,  the  edges  or  cor- 
ners of  which  have  been  broken,  is  suspended  in  a  saturated  solution 
of  the  same  substance,  it  neither  increases  nor  diminishes  in  weight. 
Yet  we  find  that  the  imperfections  are  removed,  and  that  this  takes 
place  by  the  solution  of  a  portion  of  the  substance  from  the  perfect 
surfaces  and  its  deposition  upon  the  imperfect  ones. 

Another  very  striking  proof  of  this  may  be  obtained  by  saturating 
water  with  ordinary  Glauber's  salt  (hydrated  sodium  sulphate,  Na2S04, 
10H20)  at  a  temperature  somewhat  above  the  ordinary,  say  30°.  The 
excess  of  the  solid  is  carefully  and  completely  separated  from  the 
liquid,  and  the  latter  is  allowed  to  cool  in  a  flask  loosely  stoppered  with 
cotton.  The  solution  now  contains  (Fig.  62)  a  much  larger  amount  of 


FlQ 


SOLUTION  159 

sodium  sulphate  (Ka,jS04)  than  at  its  present  temperature  it  could 
acquire  from  contact  with  Glauber's  salt.  Yet  in  the  absence  of  a 
crystal,  with  which  the  above  described  exchange  could  take  place,  no 
deposition  of  the  dissolved  substance  begins.  The  solution  may  be 
kept  indefinitely  without  alteration.  The  introduction,  however,  of  the 
minutest  fragment  of  the  decahydrate  at  once  starts  the  exchange,  and 
this  is  necessarily  very  much  to  the  disadvantage  of  the  solution  and 
the  advantage  of  the  crystal:  ~Na^04  (diss'd)  <=±  Na^SO^  lOJ^O  (solid). 
The  latter  therefore  forms  the  center  of  a  radiating  mass  of  blade- 
like  processes,  which  sprout  with  astonishing  rapidity  through  the 
liquid. 

Usually  the  cooling  of  a  concentrated  solution  leads  to  the  almost 
immediate  appearance  of  crystals  spontaneously,  and  the  substance  is 
deposited  gradually  as  the  temperature  falls.  But  solutions  of  a  nun> 
ber  of  common  substances,  such  as  sodium  thiosulphate  (photogra- 
pher's "  hypo  ")  and  sodium  chlorate,  behave  like  that  of  sodium  sulphate. 
They  are  said  to  have  a  tendency  to  give  supersaturated  solutions. 
In  general,  crystallization  can  be  started  only  by  introduction  of  a 
specimen  of  the  same  substance,  or  at  all  events  of  one  isomorphous 
(q.v.)  with  it.  The  smallest  particle  of  the  right  material  floating  in 
the  air,  if  ifc  gains  accidental  admission,  will  bring  about  the  result. 
This  shows  the  importance  of  the  interchange  of  molecules  of  which 
we  have  spoken  for  establishing  equilibrium. 

Metastdble  Condition.  —  The  above  phenomenon  is  not  an  iso- 
lated or  exceptional  one  in  physical  science.  It  is  commonly  the 
case  that,  when  the  conditions  for  some  physical  change  have  been 
reached,  the  beginning  of  the  physical  change  is  delayed  or  entirely 
fails.  The  system  is  then  said  to  be  in  a  metastable  condition.  Un- 
stable it  is  not.  Yet  it  is  not  in  the  state  of  greatest  stability,  for  the 
element  of  equilibrium  is  lacking.  Thus,  pure  water  may  easily  be 
cooled  three  or  four  degrees  below  0°  without  the  appearance  of  any 
ice.  Agitation,  however,  in  this  case,  results  in  the  appearance  of  ice 
sooner  or  later.  In  like  manner  water  may  be  heated  to  a  temperature 
above  100°  without  boiling.  Drops  of  water,  suspended  in  some  oil  of 
almost  the  same  specific  gravity,  may  even  be  raised  to  175°  before  the 
water  turns  into  steam.  Similarly,  air  which  is  saturated  with  moisture, 
if  it  contains  no  dust,  may  be  cooled  without  the  appearance  of  fog. 

Phenomena  of  supersaturation,  of  a  temporary  kind  at  least,  are  extremely 
common  in  chemistry.  Almost  every  delayed  precipitation  U  a  case  of  it.  Barium 


160  INORGANIC   CHEMISTRY 

sulphate,  for  example,  is  always  slow  in  appearing  in  dilute  solutions.  So  is  sulphur, 
set  free  from  dilute  sodium  thiosulphate  solution  by  the  action  of  an  acid.  In  the 
latter  case,  instant  reneutralization  with  a  base  does  not  prevent  the  ultimate 
appearance  of  the  sulphur,  showing  that  the  cause  does  not  lie  in  slow  interaction 
of  the  salt  with  the  acid. 

Saturation.  —  When  we  have  shaken  a  solid  for  a  sufficient  length 
of  time  with  a  given  amount  of  a  liquid,  we  obtain  a  solution  which  is 
saturated  with  respect  to  that  substance.  Having  called  this  a  satu- 
rated solution,  we  are  inclined  to  extract  from  the  term  a  meaning 
different  from  that  which  it  was  really  intended  to  convey.  We  are 
in  danger  of  thinking  that  the  solution  itself  is  in  some  way  peculiar 
—  that,  for  instance,  it  contains  all  of  the  solid  which  it  is  capable  of 
holding.  This  would  be  an  entire  misconception.  If  we  desire  to 
make  a  solution  of  sodium  sulphate  (NagSO^,  for  example,  we  may 
present  this  substance  to  the  water  either  in  the  form  of  Glauber's 
salt  (NagSC^,  10H20)  or  of  anhydrous  sodium  sulphate.  Now  the 
anhydrous  and  the  hydrated  forms  of  a  substance  always  behave  like 
entirely  different  substances.  This  hydrate  cannot  give  more  than  5 
parts  of  sodium  sulphate  (Na«jS04)  to  100  parts  of  water  at  0°.  When 
the  anhydrous  compound  is  used,  many  times  this  amount  (Fig.  62)  is 
dissolved  at  the  same  temperature  :  Na,jSO4  (solid)  <=±  NagSC^  (diss'd). 
The  solution  pressures  of  the  two  forms  are  entirely  different.  The 
phrase,  "  a  saturated  solution  of  sodium  sulphate,"  is  therefore  devoid  of 
definite  meaning.  We  must  describe  the  liquid  as  a  solution  saturated 
by  anhydrous,  or  by  hydrated  sodium  sulphate  as  the  case  may  be. 

Being  different  substances,  the  hydrated  and  anhydrous  forms  of  a  compound 
must  be  investigated  separately  as  to  their  solubility  at  various  temperatures.  The 
results  must  give  different  curves,  as  for  distinct  substances. 

Before  referring  to  the  curves  of  the  two  sodium  sulphates,  it  must  be  remarked 
that  hydrates  decompose  into  the  anhydrous,  or  some  less  hydrated  form  at  a  defi- 
nite temperature.  We  cannot  therefore  continue  the  observation  of  the  solubility 
of  the  substance  beyond  the  temperature  at  which  it  ceases  to  exist.  Thus  the 
solubility  curve  of  a  hydrate  comes  to  an  abrupt  termination  at  the  decomposition, 
or,  as  it  is  usually  called,  the  transition  point.  Now  the  decahydrate  of  sodium 
sulphate  decomposes  at  32. 4°,  so  that  its  solubilities  can  be  measured  only  from  about 
0°  to  32.4°.  The  solubility  of  the  anhydrous  form,  however,  can  be  investigated 
up  to  100°,  or  beyond  it,  if  necessary.  Not  only  so,  but  measurements  can  be 
carried  out  below  32,4°.  The  union  with  water  to  form  the  decahydrate  is  so  slow 
that  there  is  time  to  saturate  the  solution  with  the  anhydrous  body,  and  decant  the 
liquid  for  analysis,  before  the  hydrate  begins  to  be  produced. 

The  solubility  of  the  decahydrate  (Fig.  62)  rises  rapidly  between  0°  and  32.4° 
from  5  to  66  parts  in  100.  The  solubility  of  the  anhydrous  sodium  sulphate  de- 


SOLUTION  161 

creases  steadily  from  more  than  65  parts  below  32.4°  to  42.5  parts  at  100°.  The 
character  of  the  two  bodies,  in  the  matter  of  solubility,  is  therefore  entirely  differ- 
ent. The  solutions  themselves,  as  has  been  said  already,  are  identical  in  every 
way,  when  they  have  the  same  concentration,  whether  they  have  been  made  from 
the  one  substance  or  the  other. 

Let  us  return  now  to  the  proper  use  of  the  term  "  saturated  solution."  We  might 
say,  correctly,  that  at  20°  a  solution  made  from  hydrated  sodium  sulphate  and 
containing  19.4  parts  of  sodium  sulphate  in  100. of  water  was  saturated.  This 
would  not,  however,  be  the  maximum  quantity  which  the  same  amount  of 
water  could  hold,  for,  with  the  help  of  the  anhydrous  compound,  we  could  add  an 
amount  equivalent  to  prolonging  the  ordinate  at  20°  until  it  intersected  the  curve 
of  anhydrous  sodium  sulphate  somewhere  about  the  value  60.  Nor  can  we  be 
sure  that  even  then  the  water  would  contain  all  the  sodium  sulphate  which  it 
could  hold.  It  is  conceivable  that,  by  presenting  the  substance  in  some  still  other 
form,  even  greater  solubility  might  be  observed. 

A  saturated  solution,  if  we  fix  our  minds  upon  it  simply  as  a  solution,  is  not 
different  from  any  other  solution.  There  is  no  feature  in  the  properties  of  such 
a  solution  qua  solution  which  distinguishes  it  in  the  least  from  one  containing 
slightly  less  or  one  containing  slightly  more  of  the  dissolved  substance.  In  con- 
tact with  a  crystal  of  the  substance  used  to  produce  the  saturation,  however,  the 
saturated  solution  is  found  to  be  in  equilibrium,  while  the  unsaturated  solution 
takes  up  more  of  the  substance,  and  the  supersaturated  solution  at  once  deposits 
the  amount  which  it  contains  in  excess  of  the  saturated  solution.  The  words  un- 
saturated, saturated,  and  supersaturated  convey,  therefore,  no  meaning  unless  we 
add  that  the  solution  is  so  towards  some  specific  form  of  material.  These  are  quali- 
ties of  the  system  including  the  undissolved  body  and  not  of  the  solution  by  itself. 

The  fact  that  the  hydrated  and  anhydrous  sodium  sulphates  give  saturated 
solutions  of  the  same  concentration  at  32.4°,  so  that  the  curves  intersect  at  this 
point,  is  a  very  significant  one.  This  is  the  temperature  at  which  the  former  sub- 
stance turns  into  the  latter.  At  transition  points  like  this  the  values  of  solubility, 
vapor  pressure,  and  some  other  properties,  are  always  the  same  for  both  forms  (see 
Freezing-points  of  solutions). 

Properties  of  Solutions  Proportional  to  Concentration : 
Vapor  Tension.  —  Besides  osmotic  pressure  (p.  151),  there  are  several 
properties  of  solutions  which,  are  proportional  to  the  concentration  of 
the  solution. 

If,  instead  of  water,  we  introduce  aqueous  solutions  of  the  same 
substance  successively  into  the  barometric  vacuum  (Fig.  39,  p.  116),  we 
find  that  the  vapor  pressures  of  the  solutions  are  less  than  that  of 
water  at  the  same  temperature.  The  diminution  in  the  fall  of  the 
mercury  column,  which  measures  the  lowering  in  vapor  pressure,  is 
proportional  to  the  concentration  of  each  solution.  The  limit  is 
reached  with  the  saturated  solution,  although,  if  this  is  rather  con- 
centrated, the  proportionality  does  not  hold  strictly  down  to  that 
point  (see  Chap.  xvii). 


162 


INORGANIC   CHEMISTRY 


I 


Temperature 
FIG.  63. 


This  lowering  in  the  vapor  pressure  of  water  is  often  considerable. 
Thus  at  100°  a  7.5  per  cent  solution  of  potassium  chloride  shows  a 
vapor  pressure  of  only  734.1  mm.,  while  that  of  water  is  760  mm.  The 
difference  is  25.9  mm.  Hence  the  solution  has  to  be  raised  to  a  higher 
temperature  (100.96°)  before  it  boils.  This  is  almost  exactly  .037° 
per  1  mm.,  which  is  the  value  for  pure  water  (p.  118). 

This  conclusion  may  also 
be  reached  graphically  (Fig. 
63).  The  ordinates  represent 
the  vapor  tensions  correspond- 
ing to  the  temperatures  shown 
by  the  abscissae.  They  in- 
crease in  length  with  rise  in 
temperature.  The  horizontal 
dotted  line  shows  the  vapor 
tension  of  760  mm.  at  which 
any  liquid  will  boil.  The 
boiling-point  of  the  solvent  is 
therefore  the  temperature  at 
which  its  curve  intersects  this 
line.  Since  the  vapor  tensions 
of  the  solution  are  all  below 

those  of  the  solvent,  its  curve  lies  below  that  for  the  solvent  but  ascends  along  with 
the  latter.  Hence  it  also  cuts  the  760  mm.  line,  but  at  a  point  beyond  the  boil- 
ing-point of  the  pure  solvent.  Since,  for  short  lengths,  the  curves  are  very  nearly 
straight  lines,  the  distances  from  boiling-point  to  boiling-point  are  very  nearly  pro- 
portional to  the  vertical  distances  between  the  curves.  That  is  to  say,  the  eleva- 
tions in  the  boiling-point  are  proportional  to  the  depressions  in  the  vapor 
tension,  and  therefore  to  the  concentration  of  the  different  solutions  of 
any  one  substance.  When  different  substances  or  solvents  are  compared,  the 
scale  alone  is  different  (see  Chap.  xvii). 

The  same  effect  may  be  observed  in  isomorphous  mixtures  of  solid  bodies, 
which  in  many  ways  resemble  solutions.  Thus  the  vapor  tension  of  water  in  the 
alums  is  greater  than  the  average  vapor  pressure  of  water  in  the  air,  and  hence 
they  lose  their  water  of  hydration  spontaneously  (cf.  p.  121).  But  if  mixed  crystals 
of  two  alums,  say  ordinary  alum  and  iron  alum  (K2SO4,  A12(S04)3,  24H2O  and 
(NH4)2SO4,  Fe2(SO4)s,  24H2O),  are  prepared,  they  keep  perfectly.  The  vapor  tension 
of  the  water  in  each  has  been  lowered  by  the  influence  of  the  other  alum  dissolved 
in  it.  Similarly,  calcium  formate  (Ca(CH02)2,  4H20)  effloresces  (p.  123),  but  loses 
this  tendency  when  crystallized  with  some  of  the  isomorphous  barium  or  strontium 
salt. 

This  principle  explains  the  habit  of  very  soluble  compounds 
(deliquescent  substances)  to  become  moist  when  exposed  to  the  air 
and  finally  to  dissolve  in  the  water  they  seem  to  attract  from  it. 
Since,  on  account  of  their  solubility,  the  moisture  first  present  on 


SOLUTION 


163 


their  surfaces  (p.  117)  is  a  highly  concentrated  solution,  its  vapor  ten- 
sion is  lower  than  the  vapor  pressure  of  water  in  ordinary  air  (p.  161). 
Hence  further  condensation  takes  place  (pp.  121-122)  until  the  sub- 
stance is  so  diluted  that  the  vapor  tension  of  its  solution  becomes  equal 
to  the  aqueous  pressure  of  the  atmosphere.  Concentrated  sulphuric 
acid,  calcium  chloride,  and  magnesium  chloride  (p.  34)  are  substances 
of  this  kind. 


Freezing-points  of  Solutions.  —  Every  pure  liquid  has  a  defi- 
nite temperature  at  which  it  freezes.  Thus,  pure  water  freezes  at  0° 
and  benzene  at  6°.  The  presence  of  a  foreign,  dissolved  body,  how- 
ever, lowers  the  freezing-point.  Thus,  sea-water  is  harder  to  freeze 
than  fresh  water.  The  freezing-points  of  solutions  of  the  same  sub- 
stance are  found  to  be  depressed  below  that  of  the  solvent  in  pro- 
portion to  the  concentration  of  the  solute*  (see  Chap.  xvii).  This 
may  be  shown  graphi- 
cally (Fig.  64).  Theordi-  mm. 
nates  represent  vapor  fl 
tensions  corresponding  •§ 
to  the  temperatures  £ 
shown  by  the  abscissae.  | 
The  rate  at  which  the  * 

former  rise  is  greater  for      *  *  —   "-g^--- •* 

ice  than  for  water,  hence 
the  ice  curve  is  steeper. 
At  0°,  ice  and  water  can 
coexist      permanently 
(p.  115).     By  measure- 
ment,  they   have  the  same   vapor  tension  (4.6  mm.)  at   this   point. 
Theoretically,  if  they  had  not,  they  could  not  coexist    indefinitely, 
for  the  one  with  the  greater  vapor  tension  would  evaporate,  and  its 
vapor  would  condense  on  the  other  until  one  of  them  alone  remained. 

Now  the  vapor  tension  of  a  solution  is,  at  all  temperatures,  lower 
(Fig.  64)  than  that  of  water.  Hence,  for  a  solution,  the  curve  must 
cut  the  ice  curve  below  4.6  mm.  and  therefore  behind  0°.  In  other 
words,  ice  and  the  solution  cannot  have  equal  vapor  tensions,  and 
therefore  coexist  indefinitely,  except  at  some  temperature  below  0°. 

*  The  ice  which  separates  during  the  freezing  consists,  as  a  rule,  of  the  pure 
solvent,  and  the  solute  does  not  enter  into  it.  Only  when  this  is  the  case  does  t-hig 
law  represent  the  facts. 


Temp. 


0° 


FIG.  64. 


164  INORGANIC    CHEMISTRY 

But  the  temperature  at  which  ice  can  exist  indefinitely  in  a  solution 
is  the  freezing-point.  Hence,  freezing-points  of  solutions  are  always 
lower  than  those  of  the  pure  solvents. 

By  measuring  the  vapor  tensions  of  the  solution  at  several  temper- 
atures, in  order  to  see  how  far  the  curve  for  the  solution  is  below  that 
of  water,  and  then  producing  the  curve  for  the  solution  backwards,  the 
intersection  with  the  ice  curve,  and  therefore  the  freezing-point,  may  be 
obtained  graphically.  Direct  measurement  always  confirms  the  result. 

At  transition  points  like  that  of  ice  and  water,  the  values  of  proper- 
ties of  both  forms,  such  as  the  vapor  pressure  and  the  solubility  in 
some  different  substance,  are  always  identical. 

Since  the  ice  curve  is,  for  a  short  distance,  almost  a  straight  line,  it 
follows  that  the  depressions  of  the  freezing-point  are  proportional  to 
those  of  the  vapor  tension,  and  these  in  turn  are  proportional  to  the 
concentrations.  From  this  relation  we  get  the  statement  with  which 
this  section  opened.  Thus,  solutions  of  sugar  containing  11.4,  22.8,  and 
34.2.  g.  of  sugar  to  100  g.  of  water  freeze  at  -  0.62°,  - 1.23°,  and  - 1.85°, 
respectively.  Numerically,  in  the  case  of  water,  a  lowering  of  the 
vapor  pressure  by  T^7  of  its  amount  at  each  temperature  sets  the 
freezing-point  back  1.05°. 

In  everyday  life  we  scatter  salt  on  ice  to  melt  it.  The  salt  dis- 
solves in  the  moisture  on  the  surface,  and  the  ice  cannot  exist  in  pres- 
ence of  this  solution  at  0°.  It  melts,  absorbing  heat  in  doing  so,  until 
the  temperature  of  the  mixture  reaches  that  of  a  freezing,  saturated 
solution  of  salt  (about  —  21°  =  —  6°  F.).  When  the  existing  temperature 
is  lower  than  this,  the  salt  has  no  effect  on  the  ice.  Freezing  mixtures, 
being  usually  mixtures  of  ice  with  various  soluble  substances,  work 
in  accordance  with  this  principle. 

Densities  of  Solutions.  —  The  densities  or  specific  gravities  of  solutions 
are  also  functions,  although  not  simple  ones,  of  the  concentrations,  and  hence  the 
latter  are  commonly  defined  for  commercial  purposes  by  the  former.  We  purchase 
ammonium  hydrate  of  "0.88  sp.  gr.,"  meaning  35  per  cent  of  ammonia,  or  sul- 
phuric acid  of  "1.84  sp.  gr.,"  meaning  100  per  cent  of  the  acid. 

The  solution  usually  has  a  smaller  volume  than  that  of  the  sum  of  the  constit- 
uents. Thus,  58.5  g.  of  sodium  chloride  occupying  27.5  c.c.,  give,  when  dissolved 
in  10  liters  of  water,  a  solution  whose  volume  is  10.0166  liters.  The  solution  is  only 
16.6  c.c.  more  bulky  than  the  water.  Similarly,  a  formula- weight  of  potassium 
nitrate,  occupying  44.7  c.c.,  adds  only  38.5  c.c.  to  the  bulk  of  10  liters  of  water. 

Heat  of  Solution.  —  When  a  body  is  dissolved  in  water,  the  solu- 
tion may  be  either  warmer  or  colder  than  the  original  materials.  The 


SOLUTION  165 

sources  or  destiny  of  the  heat  given  out  or  absorbed  have  not  been 
studied  in  such  a  way  that  definite  statements  can  be  made  about  the 
theory  of  the  subject.  There  are  many  factors  which  would  have  to  be 
considered.  For  example,  the  body,  if  a  solid,  goes  into  an  essentially 
liquid  condition,  and  its  heat  of  fusion  is  always  negative.  The  changes 
in  its  molecular  condition  involve  either  liberation  or  absorption  of 
heat.  The  change  in  volume  must  have  a  heat  effect  connected  with 
it.  But  almost  nothing  is  known  about  these  and  other  (see  Heat 
of  ionization)  essential  parts  of  the  phenomenon. 

The  first  water  used  always  causes  a  greater  heat  change  than  the  ad- 
dition of  succeeding  equal  amounts.  Heats  of  solution  are  measured  for 
the  solution  of  one  formula-weight  of  the  substance  in  unlimited  water. 
The  values  in  calories  for  some  common  substances  are  as  follows  : 


Aq)  =  +  39,170  (NaCl,  Aq)  =  -  1180 

(HC1,  Aq)  =  +  17,400  (NH4C1,  Aq)  =  -  3880 

(KOH,  Aq)    =+12,500  (KC1,  Aq)  =  -  4440 

(NaOH,  Aq)  =  +  9780  (Na2C08,  10H20,  Aq)  =  -  16,160 

(Na^COg,  Aq)  =  -f  5640  (Na^SO^  Aq)  =  +  460 

(CaCL,,  Aq)  =  +  3258  (Na2S04,  10H2O,  Aq)  =  -  18,760 

When  a  substance  comes  out  of  solution,  the  heat  effect  is  equal  and  of 
opposite  sign  to  that  occurring  when  the  same  substance  goes  into  solu- 
tion. Hence,  since  the  decahydrate  of  sodium  sulphate  absorbs  heat 
in  dissolving,  a  considerable  development  of  heat  is  noticed  when  it 
suddenly  crystallizes  from  a  supersaturated  solution.  Some  ether  in  a 
tube  immersed  in  the  solution  may  be  boiled  by  this  heat  and  its  vapor 
set  on  fire  to  make  the  fact  evident  at  a  distance.  An  important  rela- 
tion between  heat  of  solution  and  solubility  will  be  discussed  under 
van't  Hoff's  law  of  mobile  equilibrium  (g.v.). 

Definition  of  a  Solution.  —  We  are  now  able  to  make  a  brief 
statement  which  shall  distinguish  solutions  from  mixtures  on  the  one 
hand  and  from  chemical  compounds  on  the  other.  Solutions  are 
homogeneous  mixtures  of  two  or  more  substances  which  are  not  sep- 
arable into  their  constituents  by  mechanical  means  without  altering  the 
state  of  one  of  the  substances,  and  -whose  properties  vary  continuously 
with  the  proportions  of  the  constituents  between  certain  limits. 

Application  in  Chemical  Work*  —  The  theory  of  this  subject 
has  been  given  on  account  of  its  intensely  practical  interest,  and  it 


166  INORGANIC   CHEMISTRY 

should  be  kept  in  mind  in  all  ordinary  chemical  operations.  It  will 
afford  an  explanation  of  many  things  which  might  otherwise  be  attrib- 
uted to  the  wrong  cause,  or  might  remain  entirely  without  explanation. 
For  example,  why  is  the  action  of  a  metal  upon  an  acid  so  slow  ?  We 
must  remember  that  an  acid  diluted  with  water  is  being  used,  and 
only  one  molecule  out  of  every  dozen  or  hundred  is  a  molecule  of  the 
acid.  So  that  the  access  of  the  latter  to  the  metal  is  restricted  at  first, 
and  becomes  more  and  more  so  as  the  molecules  of  the  acid  in  the 
immediate  vicinity  of  the  metal  undergo  chemical  change.*  On  the 
other  hand,  the  metal,  especially  if  it  be  in  the  form  of  sticks  obtained 
by  casting,  presents  one  of  the  elements  in  the  action  in  a  most  com- 
pact form.  The  only  parts  which  are  accessible  to  the  acid  are  those 
upon  the  surface,  and,  the  metal  not  being  appreciably  soluble  in  water, 
the  molecules  can  only  pass  off  and  expose  a  fresh  layer  very  slowly. 
It  is  no  wonder  that  many  chemical  actions  occupy  a  considerable  time. 
The  wonder  is  that  they  should  take  place  as  rapidly  as  they  do. 
Their  speed  would  seem  to  point  to  a  most  intense  chemical  activity, 
even  in  the  seemingly  feebler  instances.  Various  artifices  are  habitu- 
ally employed  for  facilitating  chemical  action.  Thus,  the  metal  may  be 
reduced  to  a  leafy  form  by  pouring  the  molten  substance  into  cold 
water.  Naturally,  with  metals,  the  maximum  surface  and  the  most 
rapid  chemical  action  are  obtained  by  using  a  fine  powder. 

The  most  speedy  interaction  of  all,  other  things  being  equal,  must 
be  attainable  by  dissolving  all  the  interacting  substances  in  water. 
Under  these  circumstances  all  the  molecules  of  each  substance  must 
simultaneously  have  many  molecules  of  the  others  within  easy  reach  of 
them. 

Exercises.  —  1.  Give  other  examples  of  limited  solubility  in  vari- 
ous solvents  (p.  146). 

2.  If  you  were  not  permitted  to  evaporate  sea-water  to  dryness,  how 
should  you  show  that  it  was  a  solution  and  not  a  pure  substance  ? 

3.  Eeexpress  Henry's  law  (p.  154)  in  terms  of  the  volume  of  gas 
dissolved  at  different  pressures. 

4.  If  hydrogen  sulphide  is  diluted  with  ten  times  its  volume  of 
hydrogen,  what  volume  of  it,  estimated  as  pure  gas,  will  be  dissolved 
by  20  volumes  of  alcohol  at  0°  and  760  mm.  (p.  154)  ? 

*  In  spite  of  the  continuous  exhaustion  of  the  acid,  there  is  often  a  steady  in- 
crease in  the  rate  at  which  a  dilute  acid  interacts  with  a  metal.  This  is  due,  at 
first,  to  the  dirt  on  the  surface  of  the  metal  which  temporarily  obstructs  the  ac- 
tion, and,  later,  to  the  rising  temperature  of  the  interacting  bodies, 


SOLUTION  167 

5.  If  the  dissolved  air,  after  being  removed  from  water  by  boiling, 
were  to  be  shaken  with  water  once  more,  in  what  proportions  by 
volume  would  the  gases  now  dissolve  (p.  155)  ? 

6.  Kead   from   the  curves   (p.  157)   the  solubilities  of  potassium 
nitrate  at  15°,  of  potassium  chloride  at  30°,  of  potassium  chlorate  at 
45°.     What  are  the  relative   rates  at  which  the  solubilities  of  these 
salts  increase  with  rise  in  temperature  ? 

7.  Express  the  concentrations  of  solutions  of  ammonium  chloride, 
saturated  at  0°  (sp.  gr.  1.076),  and  of  potassium  sulphate  K2S04,  satu- 
rated at  10°  (sp.  gr.  1.083),  in  terms  of  a  normal  solution  (p.  149). 

8.  Express  the  concentration  of  a  five  per  cent  aqueous  solution 
of  phosphoric  acid  (sp.gr.  1.027),  in  terms  of  a  normal  and  a  molar 
solution,  respectively. 

9.  Name  the  phases  (p.  156)  in  a  system  consisting  of  oxygen  and 
its  aqueous  solution,  (a)  above  0°,  (&)  below  0°. 

10.  When  a  solution  of  a  very  soluble  substance,  like  zinc  chloride, 
is  evaporated  to  dryness  on  a  water  bath,  why  is  the  escape  of  the 
last  portions  of  the  solvent  go  much  slower  than,  is  that  of  the  first  ? 


CHAPTER   XI 
CHLORINE   AND    HYDROGEN    CHLORIDE 

CHLOKINE  was  first  recognized  as  a  distinct  substance  by  Scheele 
(1774).  He  obtained  it  from  salt  by  means  of  manganese  dioxide, 
using  the  common  method  described  below.  It  was  for  years  sup- 
posed to  be  a  compound  containing  oxygen,  but  the  work  of  Davy 
(1809-1818)  established  the  fact  that  it  is  an  element. 

Occurrence.  —  Chlorine  does  not  occur  free  in  nature.  There  are, 
however,  many  compounds  of  it  to  be  found  in  the  mineral  kingdom. 
Sea-water  contains  a  number  of  chlorides  in  solution.  Nearly  2.8  of 
the  3.6  per  cent  of  solid  matter  in  sea-water  is  common  salt  (sodium 
chloride,  NaCl).  During  past  geological  ages  the  evaporation  of  sea- 
water  has  led  to  the  formation  of  immense  deposits  of  the  compounds 
usually  found  in  such  water.  Thus,  at  Stassfurt,  such  strata  attain  a 
thickness  of  over  a  thousand  feet.  Certain  layers  of  these  strata  are 
composed  mainly  of  sodium  chloride,  called  by  the  mineralogist  halite 
(rock  salt).  In  other  layers  potassium  chloride  (sylvite),  and  hydrated 
magnesium  chloride  (bischofite),  and  other  compounds  of  chlorine,  oc- 
cur. The  chloride  of  silver  (horn  silver)  is  a  valuable  ore. 

Preparation.  —  Chlorine  cannot  be  obtained  with  the  same  ease 
as  oxygen.  There  are  only  a  few  chlorides,  such  as  those  of  gold  and 
platinum,  which  lose  chlorine  when  heated,  and  they  are  too  expensive 
or  difficult  to  make  for  laboratory  use.  We  employ  therefore  methods 
like  those  used  for  the  preparation  of  hydrogen.  We  may  (1)  decom- 
pose any  chloride  by  means  of  electricity,  just  as,  to  get  hydrogen,  we 
electrolyzed  a  dilute  acid  (pp.  63,  94).  Or  (2)  we  may  take  some  inex- 
pensive compound  of  chlorine,  such  as  hydrogen  chloride  (HC1),  and  by 
means  of  some  simple  substance  which  is  capable  of  uniting  with  the 
other  constituent,  —  here  oxygen  serves  the  purpose,  —  secure  the 
liberation  of  the  element  (p.  99).  Or  (3) — and  this  turns  out  to  be 
the  most  convenient  laboratory  method  —  we  may  use  a  more  complex 
action. 

168 


CHLORINE  AND  HYDROGEN  CHLORIDE 


169 


Electrolysis  of  Chlorides.  —  Hydrogen  chloride  and  those  chlo- 
rides of  metals  which  are  soluble  in  water  are  all  decomposed  when  a 
current  of  electricity  is  passed  through  the  aqueous  solution.  They 
yield  chlorine  at  the  positive  electrode.  The  other  constituent,  the 
hydrogen  (Fig.  65),  manganese,  or  whatever  it  may  be,  is  liberated  at 
the  negative  wire.  To  decompose  hydrochloric  acid  an  electromotive 
force  of  at  least  1.31  volts  is  required.  Since  the  chlorine  is  soluble  in 
water,  the  efferves- 
cence due  to  its  re- 
lease is  not  notice- 
able until  the  liquid 
round  the  electrode 
has  become  satu- 
rated with  the  gas : 
CLj  (diss'd)  <±  C12 
(gas).  The  shape 
of  the  apparatus 
keeps  the  two  prod- 
ucts from  mingling. 
The  presence  of  the 
chlorine  in  the  liq- 
uid at  the  positive 
end  may  be  shown 
by  a  suitable  test 
(pp.  99  and  175). 

In  commerce 
chlorine  is  now  ob- 
tained chiefly  by 
this  method,  sodium 

chloride  or  potassium  chloride  being  the  source  of  the  element.  Elec- 
trodes of  artificial  graphite  are  used,  as  most  other  conductors  unite  with 
the  chlorine.  The  potassium  or  sodium,  as  the  case  may  be,  travels 
towards  the  negative  electrode,  but  is  not  liberated  (p.  99).  Instead, 
potassium  or  sodium  hydroxide  accumulates  in  the  solution  round  the 
plate  and  hydrogen  escapes.  The  chlorine  is  released  at  the  positive 
electrode,  as  usual.  The  hydroxide  and  the  chlorine  both  find  chemi- 
cal applications.  The  chlorine  is  either  liquefied  by  compression  in 
iron  cylinders  or  employed  at  once  for  making  bleaching  powder  (q.v.). 

In  the  Acker  process  (<?.#.),  melted,  impure  sodium  chloride,  with- 
out any  solvent,  is  electrolyzed. 


FIG.  65. 


170  INORGANIC   CHEMISTRY 

Action  of  Free  Oxygen  on  Chlorides.  —  Oxygen  does  not  inter- 
act with  sodium  chloride  even  at  a  high  temperature.  Hence  the 
chlorine  must  first  be  transferred  to  some  other  form  of  combination. 
By  the  interaction  of  acids,  of  which  the  most  commonly  used  is  sul- 
phuric acid,  with  chlorides,  of  which  the  cheapest  is  sodium  chloride, 
we  obtain  hydrogen  chloride.  The  details  of  this  action  are  described 
below  (p.  178).  In  order  to  liberate  chlorine  from  this  compound,  we 
may  combine  the  hydrogen  with  oxygen  obtained  from  the  air.  The 
action  is  in  accordance  with  the  equation : 

2HC1  +  0  -»  H20  +  2C1. 

The  two  gases  interact  so  slowly,  however,  that  a  catalytic  agent  must 
be  employed.  The  mixture  of  air  and  hydrogen  chloride  is  passecj 
over  pieces  of  heated  pumice-stone  or  broken  brick  previously  saturated 
with  cupric  chloride  solution.  A  temperature  of  370°-400°  is  used 
In  the  resulting  gas  the  chlorine  is  mixed  with  steam  and  with  a  very 
large  volume  of  nitrogen  which  entered  with  the  oxygen,  so  that  for 
making  the  pure  substance  this  method  (Deacon's  process)  is  quite  UD 
suitable. 

Using  the  same  principle,  magnesium  chloride  may  be  heated  in  a  stream  of  air, 
when  the  oxide  of  magnesium  is  formed  and  chlorine  is  given  off:  MgCl2  -f  O  — » 
MgO  -H  2C1.  The  oxide  of  magnesium  can  then  be  treated  with  hydrochloric 
acid  to  regenerate  the  chloride,  which  in  turn  may  be  subjected  once  more  to  th«s 
action  of  oxygen.  The  process  is  thus  a  continuous  one. 

The  above  action  is  spoken  of  as  an  oxidation.  It  is  true  that  no 
oxygen  is  actually  introduced  into  the  hydrogen  chloride  as  a  whole. 
The  removal  of  hydrogen  from  combination  with  the  chlorine  is,  how- 
ever, the  first  step  towards  the  introduction  of  oxygen  into  combination 
with  the  latter,  and  is  essentially  an  oxidation. 

Action  of  Combined  Oxygen  upon  Chlorides.  —  The  usual 
laboratory  method  of  making  chlorine  is  to  mix  a  solution  of  hydrogen 
chloride  in  water  (hydrochloric  acid,  p.  93)  with  roughly  powdered  man- 
ganese dioxide.  The  flask  containing  the  mixture  is  heated  by  means 
of  a  bath  containing  hot  water  (Fig.  66).  The  gas  is  passed  through 
a  washing  bottle  containing  water,  in  order  to  remove  some  hydro- 
gen chloride  which  may  be  carried  over.  It  may  be  dried,  if  neces- 
sary, in  a  second  washing  bottle  containing  concentrated  sulphuric 
acid.  It  cannot  be  collected  over  water  on  account  of  its  solubility,  so 
that  jars  are  usually  filled  with  it  by  downward  displacement  of  air. 


CHLORINE  AND  HYDROGEN  CHLORIDE 


171 


The  chemical  change  used  here  is  somewhat  complex.  When  an 
acid  (here  HC1)  interacts  with  an  oxide  (here  MnO2),  the  hydrogen  of 
the  former  unites  with  the  oxygen  of  the  latter,  giving  water.  We 
perceive  at  once  that  to  combine  with  20,  4H,  obtainable  only  by 
taking  4HC1,  will  be  required.  Hence  the  equation  might  be  : 


Mn02  +  4HC1  ->  2H.P  +  MnCl4. 


This   is    probably   what   happens  in  the  first   place.     The   products 
actually  obtained,  however,  are  water,  manganous  chloride  (MnCLj)  and 


FIG.  66. 


chlorine.     The  manganese  tetrachloride  is  decomposed  by  the  heating, 
the  chlorine  escapes,  and  the  other  two  products  remain  in  the  vessel. 

Mn02  +  4HC1  ->  2H20  +  MnCL,  +  2C1.  (1) 

We  owe  the  chlorine  to  the  fact  that  the  tetrachloride  is  unstable. 

When  the  mixture  is  surrounded  by  ice  and  saturated  with  chlorine,  it  can  be 
shown  with  some  degree  of  certainty  that  it  contains  the  tetrachloride.  If  it  is 
quickly  poured  into  water,  hydrated  manganese  dioxide  is  precipitated  (Wacker). 
The  decomposition  of  the  tetrachloride  is  reversible  : 

MnCl4?=i  MnCl2  +  C12, 

and  is  driven  back  by  the  excess  of  chlorine.     The  tetrachloride  is  hydrolyzed  by 
water  : 

MnCl4  +  2H20  +  xH2O  -*  MnOt,  xH2O  +  4HCL 


172  INORGANIC   CHEMISTRY 

The  action  (1)  is  of  a  type  very  common  in  chemistry.  It  is  more 
complex  even  than  double  decomposition  (p.  99),  and,  unlike  this,  its 
results  cannot  be  anticipated  by  guessing.  If  we  had  used  manganous 
oxide  (MnO),  we  should  have  had  a  double  decomposition  : 

MnO  +  2HC1  S  H20  +  MnCL,,  (2) 

but  we  should  have  got  no  chlorine.  Perhaps  the  simplest  way  to  de- 
scribe the  difference  between  these  two  actions  is  in  terms  of  the  val- 
ence of  the  manganese.  In  MnIV02H  the  element  is  quadrivalent. 
This  means  that  its  atomic  weight  professes  to  be  able  to  hold  four 
unit  weights  of  a  univalent  element.  The  four  valences  of  oxygen 
(20  n)  can  do  the  same  thing.  In  equation  (1)  the  oxygen  fulfils  this 
promise  by  taking  4H.1  But  the  MnIV  can  hold  only  2C11  permanently 
and  lets  the  other  2C11  go  free.  In  other  words,  the  valence  of  the 
unit  weight  of  manganese  changes  in  the  course  of  the  action.  In 
equation  (2),  on  the  other  hand,  the  manganese  is  bivalent  to  start 
with  (MnII011),  and  is  able  to  retain  the  amount  of  chlorine  (2C11) 
equivalent  to  0".  Actions  like  that  of  manganese  dioxide  in  (1)  are 
classed  as  oxidations.  The  hydrogen  chloride,  or  rather  half  of  it,  is 
oxidized.  A  graphic  mode  of  writing  may  make  this  remark  clearer  : 

2HC1->H20  +  MnnCl2 


The  upper  half  is  a  double  decomposition,  the  lower  an  oxidation  by 
half  the  combined  oxygen  of  the  dioxide. 

In  practice,  instead  of  employing  aqueous  hydrochloric  acid  we 
frequently  use  the  materials  from  which  it  is  prepared,  namely,  common 
salt  and  concentrated  sulphuric  acid  (p.  178),  along  with  the  manganese 
dioxide.  Under  those  circumstances,  the  action  appears  more  com- 
-plex,  but  is  simply  a  combination  of  the  two  chemical  changes,  and  is 
represented  by  the  equation  : 
Mn02  +  2NaCl  +  3H2S04  ->  2H20  +  2NaHS04  +  MnS04  +  2C1. 

Using  the  same  principle,  we  find  that  lead  dioxide,  sodium  chlorate, 
potassium  dichromate  (K2O2O7),  potassium  permanganate  (KMn04), 
and  many  other  substances,  when  treated  with  aqueous  hydrochloric 
acid,  likewise  give  chlorine. 

Perhaps  the  best  way  to  obtain  a  steady,  easily  regulated  stream  of 
chlorine  is  to  place  some  solid  potassium  permanganate  in  a  flask, 
arranged  like  that  in  Fig.  70  (p.  232)  but  without  the  U  -tubes,  and 


CHLORINE  AND   HYDROGEN   CHLORIDE  173 

allow  concentrated  commercial  hydrochloric  acid,  diluted  with  half  its 
volume  of  water,  to  fall  upon  it  drop  by  drop  from  the  dropping  fun- 
nel. The  action  is  very  rapid,  the  acid  is  exhausted  almost  as  fast  as 
it  falls,  and  so  the  stream  of  gas  can  be  stopped  by  simply  closing  the 
stopcock.  In  this,  and  in  all  the  cases  mentioned  above,  the  chemical 
action  follows  the  same  plan  as  before  (p.  171).  The  oxygen  and  hy- 
drogen combine  to  form  water,  hence  here  8HC1  will  be  needed ;  the 
metals  yield  the  chlorides  which  are  stable  at  the  temperature  of  the 
action,  and  the  rest  of  the  chlorine  is  liberated  : 

KMn04  +  8HC1  ->  4H2O  f  KC1  -f  MnCL,  -{-  5C1. 

Kinetic-Molecular  Hypothesis  Applied  to  these  Actions.  —  In 

preparing  chlorine  by  the  usual  laboratory  method  (p.  170),  it  will  be 
observed  that  the  gas  is  produced  rather  slowly.  Even  heating  does 
not  accelerate  the  interaction  very  greatly.  The  situation  is  that  we 
have  placed  together  manganese  dioxide  in  a  granular  form  and  water 
which  contains  hydrogen  chloride  in  solution.  The  dioxide  is  very 
insoluble  in  water,  and  consequently  its  molecules,  which  must  dissolve 
before  they  can  meet  the  acid,  become  available  very  slowly :  MnO2 
(solid)  <=±  Mn02  (diss'd).  The  finer  the  pulverization,  the  less  will  be 
the  delay  from  this  cause.  On  the  other  hand,  the  acid  contains 
originally  only  about  one  molecule  of  hydrogen  chloride  for  every  four 
of  water,  and  as  the  former  is  used  up  the  scarcity  of  the  active  sub- 
stance becomes  greater. 

Again,  we  heat  the  mixture  on  a  water  bath  so  as  to  hasten  the 
process  (p.  72)  by  raising  the  temperature  to  about  90°.  When  we 
prepared  oxygen,  we  forced  the  temperature  up  with  a  naked  Bunsen 
flame  until,  between  200°  and  300°,  a  sufficiently  rapid  stream  of  the 
gas  was  secured.  The  iron  and  sulphur  (p.  11)  we  raised  nearly  to  a 
red  heat.  Here  the  conditions  make  stronger  heating  impossible.  No 
aqueous  solution  of  hydrogen  chloride  can  be  raised  above  110°,  the 
maximum  boiling-point  (p.  182).  But  we  must  not  carry  the  heating 
so  far  as  110°,  because  even  below  this  point  the  concentrated  acid  gives 
off  gaseous  hydrogen  chloride  freely.  If  we  did,  we  should  contaminate 
our  chlorine  and  at  the  same  time  lose  a  part  of  one  of  the  ingredients  on 
which  the  action  depends.  Intelligent  chemical  work  always  demands 
a  careful  consideration  of  purely  physical  facts  of  this  description. 

Physical  Properties.  —  Chlorine  differs  from  the  gases  we  have 
encountered  so  far  in  having  a  strong  greenish-yellow  tint  (Gk. 


174  INOEGANIC   CHEMISTRY 


o's,  pale  green),  a  fact  which  gave  rise  to  its  name,  and  having  a 
powerful  irritating  effect  upon  the  membranes  of  the  nose  and  throat. 

Density  (H  =  1),  35.79  Boiling-point  (liq.),  -  33.6° 

Weight  of  1  1.,  3.220  g.  Melting-point  (solid),  -  102° 

Solubility  in  Aq.  (20°),  216  vols.  in  100  Vap.  tension  (liq.)  0°,  3.66  atmos. 

Crit.  temp.,  +  146°  Vap.  tension  (liq.)  20°,  6.62  atmos. 

Since  a  liter  of  air  weighs  1.293  g.,  chlorine  is  two  and  a  half 
times  heavier.  In  solubility  it  stands  between  slightly  soluble  gases, 
like  oxygen  and  hydrogen,  and  those  which  are  extremely  soluble.  It 
can  be  collected  over  hot  water  or  a  strong  solution  of  salt. 

It  was  first  liquefied  by  Northmore.  The  critical  temperature  (p. 
133)  is  exceptionally  high  (146°),  so  that  at  all  ordinary  temperatures 
the  gas  can  be  liquefied  by  compression  alone.  It  forms  a  yellow 
liquid  which,  contained  in  steel  cylinders,  is  now  an  article  of  com- 
merce. On  being  cooled  below  —  102°,  it  gives  a  pale-yellow  solid. 

Chemical  Properties.  —  Chlorine  is  at  least  as  active  a  sub- 
stance as  is  oxygen.  It  presents  a  more  varied  array  of  chemical 
properties  than  does  that  element. 

1.  When  a  strong  solution  of  chlorine  in  water  is  cooled  with  ice, 
crystals  of  chlorine  hydrate  (Cl,  4H20)   are  formed.     This  hydrate, 
being  the  hydrate  of  a  gas,  has  a  high  vapor  tension  of  chlorine,  as 
well  as  a  smaller  one  of  water  (p.  121).     At  0°  the  partial  tension  of 
the  chlorine  is  249  mm.,  at  9.6°  it  is  760  mm.     The  compound  decom- 
poses rapidly,  therefore,  at  the  latter  temperature  unless  tlie  pressure 
of  chlorine  over  it  is  greater  than  one  atmosphere. 

2.  Chlorine  unites  directly  with  many  elements.    A  jet  of  hydrogen 
burns  vigorously   in    chlorine,    producing    hydrogen    chloride.     The 
presence  of  this  product  may  be  recognized  at  once,  because,  while 
chlorine    in    contact   with   moist    breath   gives   no   cloud,   hydrogen 
chloride  (q.v.)  produces  a  dense  fog.     The  union  of  the  gases,  when 
a  mixture  of  them  is  kept  cold  and  in  the  dark,  is  too  slow  to  be  per- 
ceived.    On  exposure  to  diffused  light,  however,  they  unite  slowly, 
while  a  sudden  flash  of  sunlight  or  the  burning  of  a  magnesium  ribbon 
causes  instant  explosion.     The  function  of  the  light  here  is  entirely 
different  from  that  in  the  decomposition  of  silver   chloride   (p.  19). 
In  the  latter  case  light  was  used  to  maintain  the  change,  which  comes  to 
a  stop  whenever  the  light  is  withdrawn.     The  action  was  endothermal 
and  consumed  energy.     The  union  of  hydrogen  and  chlorine  is  highly 
exothermal,  and  a  minimum  of  light  only  is  needed  to  start  it  (p.  74). 


CHLORINE  AND   HYDROGEN  CHLORIDE  175 

Sodium  burns  in  chlorine,  producing  a  cloud  of  white  particles  of 
sodium  chloride.  Copper  in  the  condition  of  thin  leaf  commonly 
used  for  gilding  (Dutch-metal),  catches  fire  spontaneously  when  thrust 
into  the  gas.  Phosphorus  burns  in  it  with  a  rather  feeble  light,  pro- 
ducing phosphorus  trichloride,  a  liquid  (b.-p.  74°)  which  condenses 
readily.  The  proportions  of  the  materials  taking  part  in  this  change 
show  that  one  atomic  weight  of  phosphorus  (31)  and  three  atomic 
weights  of  chlorine  form  the  product : 

p  +  3d  ->  PCI,. 

Almost  all  the  more  familiar  elements  unite  directly  with  chlorine  to 
form  chlorides.  The  exceptions  are,  nitrogen,  oxygen,  carbon,  helium, 
and  argon  (p.  68).  Compounds  with  all  but  the  last  two  may  be  ob- 
tained as  products  of  more  complex  interactions,  however. 

Some  elements  form  more  than  one  chloride.  Thus,  when  phos- 
phorus unites  at  a  high  temperature  with  chlorine,  we  get  phosphorus 
trichloride.  When  we  pass  chlorine  gas  into  cooled  phosphorus  tri- 
chloride, however,  a  considerable  absorption  takes  place,  and  finally 
a  solid  body  containing  two  more  atomic  weights  of  chlorine,  phos- 
phorus pentachloride  (PC15),  is  formed. 

Carefully  dried  substances  unite  slowly  or,  to  all  appearance,  not 
at  all  with  chlorine.  Thus,  perfectly  dry  chlorine  does  not  unite  with 
copper,  even  when  the  system  is  warmed.  The  introduction  of  a  drop 
of  water  (p.  75),  however,  into  a  remote  part  of  the  apparatus  at 
once  supplies  the  trace  of  moisture  which  seems  to  be  necessary  to 
start  the  chemical  change.  Chlorine,  as  ordinarily  made,  unites  with 
iron  with  great  vigor.  But  dried  chlorine  is  quite  indifferent  and 
does  not  attack  the  steel  cylinders  in  which  it  is  now  sold.  The  water 
is  a  catalytic  agent,  and  we  regard  it  as  simply  hastening  an  action 
which  otherwise  is  vanishingly  slow  (p.  73). 

3.  Chlorine,  like  sodium  (p.  97),  may  also  displace  elements  which 
are  already  in  combination.  Thus,  when  turpentine  (C10H16)  is  poured 
over  a  strip  of  paper  and  is  then  immersed  in  a  jar  of  chlorine,  a 
violent  action  takes  place,  and  an  immense  cloud  of  finely  divided 
carbon  bursts  forth.  The  heat  of  the  action,  as  it  starts,  vaporizes 
the  turpentine,  and  the  hydrogen  in  the  latter  unites  with  the  chlorine 
forming  hydrogen  chloride,  while  the  carbon  is  set  free :  (C10HM  + 
16C1  _>  16HC1  +  IOC). 

The  action  of  chlorine  on  potassium  iodide,  dry  or  in  solution,  is  of 
this  kind,  and  furnishes  the  commonest  test  (p.  99)  for  free  chlorine. 


176  INORGANIC  CHEMISTRY 

The  chlorine  simply  takes  the  place  of  the  iodine  (KI  4  Cl  — >  KC1  4 1), 
and  the  latter  is  liberated.  Iodine  when  moist  is  deep  brown  in  color, 
and  the  amount  liberated  by  a  large  quantity  of  chlorine  may  easily 
be  seen.  Yet,  it  is  an  advantage  so  to  arrange  a  test  that  it  may  be  as 
delicate  as  possible,  that  is,  may  give  a  plainly  visible  result  with  the 
minimum  of  material.  In  this  case  much  starch  emulsion  and  a  little 
potassium  iodide  are  employed.  Strips  of  filter  paper  dipped  in  this 
mixture  show  a  deep-blue  color  (see  Iodine)  when  brought  into  a  gas 
containing  even  a  trace  of  free  chlorine.  Combined  chlorine,  as  in  a 
chloride,  has  no  effect. 

4.  When  actions  like  the  above  are  moderated  by  proper  means, 
the  decomposition  is  not  so  complete.  If  methane  (marsh  gas,  CH4)  is 
mixed  with  chlorine  and  exposed  to  sunlight,  a  slower  action  occurs, 
of  which  the  first  stage  consists  in  the  removal  of  one  unit  weight  of 
hydrogen  and  the  substitution  of  chlorine  for  it  according  to  the 
following  equation : 

CH4  +  2C1  ->  CH8C1  +  HC1. 

The  process  may  continue  further  by  the  substitution*  of  chlorine  for 
the  units  of  hydrogen  one  by  one  until  carbon  tetrachloride  (CC14)  is 
finally  formed. 

A  most  interesting  and  important  action  of  this  class  occurs  when 
chlorine  is  dissolved  in  water.  A  small  part  of  the  chlorine  interacts 
with  a  little  of  the  water,  the  element  being  substituted  for  one-half 
of  the  hydrogen  : 

2C1  4-  HHO  ->  HOC1  4  HC1.  (1) 

Only  traces  of  hydrogen  chloride  and  hypochlorous  acid  are  produced, 
however.  The  change  comes  quickly  to  a  standstill,  because  the  prod- 
ucts interact  even  more  vigorously  to  reproduce  chlorine  and  water : 

HC1  4-  HOC1  <->  2C1  4  H20,  (2) 

the  interaction  being  reversible  (p.  64).  This  interaction  of  chlorine 
and  water  (1),  slight  as  it  is,  is  of  importance  on  account  of  the  insta- 

*  Substitution  resembles  displacement  (p.  99)  in  that  an  element  and  a  com- 
pound interact,  and  the  element  takes  the  place  of  one  unit  in  the  composition  of 
the  latter.  In  the  above  action,  one  unit  of  chlorine  takes  the  place  of  one  unit 
of  hydrogen.  But  the  latter  is  not  liberated  :  it  combines  with  another  unit  of 
chlorine.  Double  decomposition  the  action  is  not,  because  elements  do  not  decom. 
pose.  The  name  used  is  intended  to  fix  the  attention  on  the  compound  and  on  the 
fact  that  one  unit  has  been  substituted  for  another  in  it.  This  conception  is  a 
favorite  one  in  the  chemistry  of  compounds  of  carbon. 


CHLORINE   AND   HYDROGEN   CHLORIDE  177 

bility  of  the  hypochlorous  acid  (g>v.)  which  it  produces.  When  the 
solution  is  exposed  to  sunlight,  the  hypochlorous  acid  decomposes  and 
oxygen  gas  is  produced  :  HC10  — >  HC1  +  0.  Since  this  removes  the 
substance  on  whose  interaction  with  the  hydrogen  chloride  in  (2),  the 
reversal  of  (1)  depends,  the  latter  action  proceeds  under  continuous 
illumination  gradually  to  completion.  Hence  the  aqueous  solution  of 
chlorine  must  be  kept  in  the  dark,  since  otherwise  a  dilute  solution 
of  hydrogen  chloride  alone  remains. 

The  so-called  bleaching  action  of  "  chlorine  "  is  almost  always  the  result  of 
oxidation  of  the  coloring  matter  by  hypochlorous  acid.  Chlorine  and  the  dye  in 
the  cloth,  even  when  only  moderately  dry,  show  no  tendency  to  interaction.  This 
may  be  demonstrated  by  collecting  some  chlorine  in  a  stoppered  bottle  in  the 
bottom  of  which  a  little  concentrated  sulphuric  acid  stands.  A  piece  of  colored 
calico  may  be  attached  by  a  pin  to  a  cork  stuck  in  the  bottom  of  the  stopper  and 
so  suspended  in  the  gas.  After  twenty-four  hours  no  action  will  be  found  to  have 
occurred.  Yet  if  the  rag  is  first  moistened,  the  bleaching  is  almost  instantaneous. 
Hence  this  bleaching  action  is  treated  under  the  properties  of  hypochlorous  acid. 

5.  Chlorine  may  simply  add  itself  to  a  compound.  Thus,  one  of  the 
oxides  of  carbon,  carbon  monoxide  (CO),  when  mixed  with  chlorine 
and  exposed  to  sunlight  gives  drops  of  a  volatile  liquid  (b.-p.  8.2°) 
known  as  phosgene  (COC12). 

Chemical  Relations  of  the  Element.*  —  In  the  formation  of 
chlorides,  an  atomic  weight  of  chlorine  is  equivalent  to  one  atomic 
weight  of  hydrogen  or  of  sodium.  The  element  is,  therefore,  univalent 
(p.  103).  It  never  shows  any  higher  valence  than  this  save  in  its  oxygen 
compounds  (see  Chap.  xvi).  The  oxides  of  chlorine  interact  with 
water  to  give  acids,  and  the  element  is,  therefore,  to  be  classed  as  a 
non-metal  (p.  119).  It  belongs  to  that  group  of  the  non-metals  called 
the  halogens  (y.v.),  as  a  consideration  of  some  others  of  its  relations 
will  show  (see  Chap.  xiv). 

Uses  of  Chlorine.  —  Large  quantities  of  chlorine  are  manu- 
factured for  the  preparation  of  bleaching  materials  and  disinfecting 
agents. 

*  In  accordance  with  the  distinction  that  must  be  drawn  (p.  32)  between  the  ele- 
ment as  a  variety  of  matter  in  combination,  and  the  elementary  substance  or  free 
form  of  the  element,  and  to  avoid  a  common  source  of  confusion,  we  shall  always 
give  only  the  behavior  of  the  elementary  substance  under  the  title  chemical  proper- 
ties. The  characteristics  which  distinguish  the  compounds  of  the  element,  as  a 
class,  from,  or  relate  them  as  a  class  to  the  compounds  of  other  elements  will  then 
appear  in  a  separate  section  under  the  above  title  (see  Chap,  xiv,  first  section). 


178  INORGANIC   CHEMISTRY 

In  disinfection,  the  minute  germs  of  disease  and  putrefaction  are 
acted  upon  either  by  the  chlorine  or  by  the  hypochlorous  acid  formed 
by  its  interaction  with  water,  and  instantly  their  life  is  destroyed. 
One  of  the  processes  for  the  extraction  of  gold  involves  an  action  of 
chlorine  gas  upon  the  material  after  it  has  received  preliminary  treat- 
ment. A  chloride  of  gold  is  formed,  which  can  be  dissolved  out  of 
the  matrix  by  means  of  water,  and  the  metallic  gold  is  afterwards 
precipitated  from  the  solution. 

HYDROGEN  CHLORIDE. 

Preparation  from  Sodium  Chloride*  —  As  we  have  seen,  the 
direct  union  of  hydrogen  and  chlorine  produces  a  gas,  hydrogen  chlo- 
ride (HC1).  For  the  purpose  of  preparing  this  gas,  however,  some 
more  easily  managed  method  will  naturally  be  employed. 

In  commerce  large  quantities  of  hydrogen  chloride  are  obtained  as 
a  by-product  in  connection  with  the  manufacture  of  soda.  The  same 
materials  are  commonly  employed  in  the  laboratory.  Common  salt  is 
treated  with  concentrated  sulphuric  acid  at  a  gentle  heat.  Efferves- 
cence is  seen  and  hydrogen  chloride  is  given  off,  while  a  compound 
known  as  sodium  hydrogen  sulphate  (sodium  bisulphate)  remains 
behind.  The  action  is  represented  by  the  equation : 

NaCl  +  H2S04  ->  HC1  f  +  NaHSO,.  (1) 

The  apparatus  used  in  the  preparation  of  chlorine  (Fig.  66)  may  be 
employed.  On  account  of  its  extreme  solubility  the  gas  is  not  washed 
in  water,  however.  For  the  same  reason  it  must  be  collected  by  down- 
ward displacement  of  air,  or  over  mercury. 

The  above  statements  apply  to  the  action  of  a  large  amount  of  sul- 
phuric acid  upon  a  limited  amount  of  salt,  where  no  high  temperature 
is  employed.  If,  however,  a  larger  proportion  of  salt  is  used  and  a 
sufficiently  high  temperature  produced  by  artificial  heating,  then  com- 
mon sodium  sulphate  is  formed  according  to  the  equation : 

2NaCl  +  H2S04  ->  2HC1  +  Na2S04.  (2) 

The  former  action  is  that  which  occurs  under  the  conditions  used  in 
the  laboratory.  The  latter  is  the  action  which  is  employed  in  com- 
merce, since  it  is  for  the  purpose  of  making  sodium  sulphate  (Na2S04), 
from  which  sodium  carbonate  is  afterwards  to  be  prepared,  that  the 
operation  is  undertaken.  The  hydrogen  chloride  passes  through  a 
tower,  down  which  water  trickles  over  lumps  of  coke,  and  is  dissolved. 


CHLORINE  AND  HYDROGEN   CHLORIDE  179 

The  aqueous  solution  is  called  hydrochloric  acid,  or,  in  commerce, 
muriatic  acid. 

Interaction  of  Acids  and  Chlorides.  —  It  should  be  noted  that 
the  above  is  simply  an  illustration  of  a  perfectly  general  method. 
Almost  any  chloride  of  a  metal  might  have  been  used  instead  of  sodium 
chloride,  and  almost  any  acid  which  could  be  obtained  free  from  any 
large  amount  of  water  (see  below)  could  have  taken  the  place  of  sul- 
phuric acid.  Thus,  concentrated  phosphoric  acid  with  any  chloride 
will  give  a  change  parallel  to  the  above.  With  sodium  chloride  this 
action  would  be  NaCl  +  H8PO4  -> HC1  \  +  NaH2P04  (primary  sodium 
phosphate). 

If  various  chlorides  are  used  with  the  same  acid,  it  will  be  found 
that  the  vigor  of  the  actions  is  very  different.  In  some,  hydrogen 
chloride  will  be  produced  copiously  without  the  assistance  of  heat. 
In  others,  there  will  be  difficulty  in  showing  that  hydrogen  chloride 
gas  is  produced  at  all.  We  must  not  hastily  assume  that  this  is  owing 
to  any  greater  chemical  affinity  in  one  case  than  another.  More  ex- 
tensive experimentation  will  show  that  the  more  soluble  chlorides  as  a 
rule  give  more  vigorous  effects  than  those  which  are  less  so  (pp.  166, 173). 
Ammonium  chloride  with  sulphuric  acid  would  represent  the  former 
variety,  while  mercuric  chloride  with  the  same  acid  would  represent 
the  latter. 

The  Kinetic  Hypothesis  Applied  to  the  Interaction  of  Sul- 
phuric Acid  and  Salt.  —  One  who  has  used  the  above  method  for  mak- 
ing hydrogen  chloride  without  reflection  would  not  realize  the  complex- 
ity of  the  machinery  by  which  the  result  is  achieved.  The  means  are 
apparently  very  simple.  Yet  the  mechanical  features  of  this  experi- 
ment, when  laid  bare,  are  extremely  curious  and  interesting.  A  single 
fact  will  show  the  possibilities  which  are  concealed  in  it. 

If  we  take  a  saturated  solution  of  sodium  hydrogen  sulphate  in 
water  and  add  to  it  a  concentrated  solution  of  hydrogen  chloride  in 
water  (concentrated  hydrochloric  acid),  we  shall  perceive  at  once  .the 
formation  of  a  copious  precipitate.  This  is  composed  entirely  of 
minute  cubes  of  sodium  chloride : 

NaHS04  +  HC1  ->  H2S04  +  NaClj.  (3) 

Now  this  action  is  nothing  less  than  the  precise  reverse  of  (1),  yet  it 
proceeds  with  equal  success.  In  fact,  this  chemical  interaction  is  not 
only  reversible  (p.  64),  but  can  be  carried  to  completion  in  either  direc- 


180  INORGANIC   CHEMISTRY 

tion.  It  is  only  in  presence  of  a  large  amount  of  water  that  it  stops 
midway  in  its  career  and  is  valueless  for  securing  a  complete  trans- 
formation in  either  direction  : 

NaHS04  +  HC1  <r>  H2S04  +  NaCl. 

In  an  action  which  is  reversible,  if  the  products  remain  as  perfectly 
mixed  and  accessible  to  each  other  as  were  the  initial  substances,  their 
interaction  will  continually  undo  a  part  of  the  work  of  the  forward 
direction  of  the  change.  Hence,  in  such  a  case  the  reaction  must,  and 
does,  come  to  a  standstill  while  as  yet  only  partly  accomplished 
(cf.  p.  176)  ;  but  this  was  not  the  case  with  actions  (1)  and  (3).  Let 
us  examine  the  means  by  which  the  premature  cessation  of  each  was 
avoided. 

In  (1)  the  salt  dissolved  to  some  extent  in  the  sulphuric  acid, 
NaCl  (solid)  <z±  NaCl  (diss'd),  and  so,  by  contact  of  the  two  kinds  of 
molecules,  the  products  were  formed.  On  the  other  hand,  the  hydrogen 
chloride,  being  insoluble  in  sulphuric  acid,  escaped  as  fast  as  it  was 
formed :  HC1  (diss'd)  ^±  HC1  (gas).  Hence,  in  that  case,  almost  no 
reverse  action  was  possible,  and  the  double  decomposition  went  on  to 
completion.  With  all  the  sodium  hydrogen  sulphate  in  the  bottom  of 
the  flask,  and  most  of  the  hydrogen  chloride  in  the  space  above,  the  two 
products  might  as  well  have  been  in  separate  vessels  so  far  as  any 
efficient  interaction  was  concerned.  This  plan,  in  which  water  is  pur- 
posely excluded,  forms  therefore  the  method  of  making  hydrogen 
chloride. 

In  (3),  on  the  other  hand,  the  hydrogen  chloride  was  taken  in 
aqueous  solution,  and  was  kept  permanently  in  full  contact  with  the  so- 
dium bisulphate.  It  had  therefore  in  this  case  every  opportunity  to 
interact  with  the  latter  and  no  chance  of  escape.  Every  molecule  of 
each  ingredient  could  reach  every  molecule  of  the  other  with  equal 
ease.  Furthermore,  the  sodium  chloride  produced  as  a  result  of  their 
activity  is  not  very  soluble  in  concentrated  hydrochloric  acid  (far 
less  so  than  in  water),  and  so  it  came  out  as  a  precipitate :  NaCl 
(diss'd)  <=±  NaCl  (solid).  But  this  was  almost  the  same  as  if  it  had 
gone  off  as  a  gas.  It  meant  that  the  greater  part  of  the  salt  was  in  the 
solid  form.  It  was  in  a  state  of  fine  powder,  it  is  true.  But,  in  the 
molecular  point  of  view,  the  smallest  particle  of  a  powder  contains 
millions  of  molecules,  and  most  of  these  are  necessarily  buried  in  the 
interior.  Thus  the  sodium  chloride  was  no  longer  able  to  interact 
effectively  molecule  to  molecule  with  the  other  product,  the  sulphuric 


CHLORINE   AND  HYDROGEN  CHLORIDE  181 

acid.  Hence,  there  was  little  reverse  action  to  impede  the  progress  of 
the  primary  one.  Thus  (3)  is  nearly  as  perfect  a  way  of  liberating  sul- 
phuric acid  as  (1)  is  of  liberating  hydrogen  chloride. 

This  discussion  is  given  to  show  that,  in  many  chemical  actions, 
the  affinity  is  entirely  subordinated  by  the  effects  of  a  purely  mechan- 
ical arrangement  (cf.  pp.  28,  110,  166).  If  the  latter  is  well  devised, 
an  action  propelled  by  a  feeble  affinity  may  prevail  against  a  reverse 
action  involving  a  very  powerful  one  (see  Chemical  equilibrium). 

The  egregious  misconception  that  sulphuric  acid  is  shown  by  this  action  to  be 
"  stronger  "  than  hydrochloric  acid  was  disposed  of,  so  far  as  the  science  was  con- 
cerned, half  a  century  ago.  But  it  survives  in  suburban  chemical  circles  with  re- 
markable tenacity.  The  fact,  quaintly  enough,  is  that  the  real  relation  in  respect 
to  activity  is  just  the  reverse. 

Other  Ways  of  Obtaining  Hydrogen  Chloride.  —  Although 
never  used  for  generating  hydrogen  chloride  on  a  large  scale,  there  is 
another  important  kind  of  action  in  which  the  substance  is  a  product. 
When  water  acts  upon  the  chlorides  of  non-metallic  substances  like 
sulphur,  phosphorus,  and  iodine,  a  double  decomposition  occurs.  Since 
water  is  always  one  of  the  interacting  substances,  this  kind  of  change, 
—  a  double  decomposition  involving  -water,  —  is  called  hydrolysis 
(G-k.  v8<op,  water,  and  AiW,  the  act  of  loosing).  Thus,  when  a  little 
water  is  added  to  one  of  the  chlorides  of  phosphorus,  hydrogen  chloride 
is  formed.  Besides  this,  the  trichloride  gives  phosphorous  acid,' and 
the  pentachloride,  phosphoric  acid  : 

PC18  +  3HOH  ->  3HC1  +  P(OH)3, 
PC16  +  4H2O  ->  5HC1  +  H8P04. 

A  dissociation  is  a  reversible  decomposition  of  one  substance  into  two.  Hy- 
drolysis is  an  ordinary  double  decomposition  or  metathesis  where  water  is  one  of 
the  reagents.  Yet  it  has  been  perversely  named  hydrolytic  dissociation  by 
many  writers.  A  whole  chapter  might  be  devoted  to  the  ingenuity  with  which 
chemists  have  misnamed  many  of  the  things  with  which  they  deal.  Perhaps  this 
tendency  is  a  survival  of  the  habit  the  chemists  had  of  using  obscure  and  symbolical 
names  for  their  materials  to  prevent  the  penetration  of  their  secrets  by  uninitiated 
seekers  after  knowledge.  Important  facts  and  principles  have  been  sedulously 
labeled  with  misleading  titles,  like  :  Water  of  crystallization,  which  has  no  more  to 
do  with  crystallization  than  with  color,  density,  or  any  other  physical  property  ; 
supersaturated  solution,  which,  as  a  solution,  is  the  same  as  any  other  ;  mass  ac- 
tion, which  has  nothing  to  do  with  mass,  but  is  concerned  wholly  with  concentra- 
tion ;  strong  acid,  which  refers  to  activity  and  not  power  of  resistance  ;  reciprocal 
proportions,  a  law  in  which  reciprocals  of  numbers  play  no  part ;  downward  dis- 
placement of  air,  when  the  air  is  displaced  upwards,  and  so  forth.  Here  there 


182  INORGANIC   CHEMISTRY 

is  an  opportunity  to  confuse  hydrolysis  with  electrolytic  dissociation,  and  the  be- 
ginner never  fails  to  embrace  it.  Hydrolytic  double  decomposition  would  have 
been  a  correct  if  somewhat  clumsy  term. 

Often,  when  a  steady  stream  of  hydrogen  chloride  is  required,  con- 
centrated  hydrochloric  acid  is  placed  in  a  generating  flask,  and  concen- 
trated sulphuric  acid  is  allowed  to  trickle  into  it  from  a  dropping 
funnel.  The  hydrogen  chloride  is  less  soluble  in  diluted  sulphuric 
acid  than  in  water  (see  Product  of  solubility)  and  escapes. 

Physical  Properties.  —  Hydrogen  chloride  is  a  colorless  gas, 
which  produces  a  suffocating  effect  when  breathed. 

Density  (H  =  1),  18.23  Crit.  temp.,  +  52° 

Weight  of  11.,  1.641  g.  Boiling-point  (liq.),  -  83.7° 

Solubility  in  Aq.  (0°),  60,300  vols.  in  100      Melting-point  (solid),  -  110° 

The  gas  is  one-fourth  heavier  than  air.  On  account  of  its  great 
solubility  and  the  small  vapor  tension  of  its  solution,  it  condenses  at- 
mospheric moisture  into  a  fog  of  drops  of  hydrochloric  acid.  On 
account  of  its  high  critical  point,  it  may  be  liquefied  by  pressure  alone. 
Both  in  the  gaseous  and  liquefied  states  it  is  a  nonconductor  of  elec- 
tricity. Its  heat  of  solution  (p.  164)  is  17,400  calories. 

On  account  of  its  high  concentration,  the  solution  may  be  looked 
upon  as  a  mixture  of  liquefied  hydrogen  chloride  and  water.  At  15° 
and  760  mm.  454.6  volumes  of  the  gas  dissolve  in  1  volume  of  water, 
or  746  g.  in  1  1.  The  mixture  weighs  therefore  1746  g.  (42.7  per 
cent  of  HC1).  Its  sp.  gr.  is  1.215.  The  volume  of  the  solution  is 
given  by  the  proportion  1215  :  1  :  :  1746  :  x,  in  which  x  —  1.437  1. 
Hence  the  addition  of  454.6  liters  of  the  gas  has  increased  the  volume 
by  only  437  c.c.  Now  at  15°  the  sp.  gr.  of  liquefied  hydrogen  chlo- 
ride is  0.8320,  and  the  volume  of  746  g.  is  therefore  746  •*•  0.832  = 
896  c.c.  So  that  even  if  the  substances  had  been  mixed  in  liquid 
form  a  considerable  shrinkage  would  still  have  occurred. 

When  the  concentrated  aqueous  solution  is  heated,  it  is  the  gas  and 
not  the  water  which  is  driven  out  for  the  most  part.  When  the  con- 
centration has  been  reduced  to  20.2  per  cent  the  rest  of  the  mixture 
distils  unchanged  at  110°.  If  a  dilute  solution  is  used,  water  is  the 
chief  product  of  distillation  (about  100°),  but  gradually  the  boiling- 
point  rises,  and,  when  the  concentration  has  reached  20.2  per  cent  once 
more,  the  same  hydrochloric  acid  of  constant  boiling-point,  as  it  is 
called,  forms  the  residue.  It  is  thus  impossible  to  separate  by  distil- 


CHLORINE   AND   HYDROGEN  CHLORIDE  183 

lation  the  components  of  mixtures  which  behave  in  this  way.  This 
must  necessarily  be  the  case  whenever,  as  here,  the  vapor  tensions  of 
the  components  separately  and  those  of  all  other  mixtures  are  higher 
than  that  of  one  particular  mixture.  When,  as  is  more  often  the  case, 
one  of  the  components  has  a  vapor  tension  which  is  lower  than  that  of 
the  other  and  lower  than  that  of  any  mixture  of  the  two,  this  compo- 
nent will  tend  to  remain  behind,  and  separation  can  be  effected.  The 
separation  of  petroleum  products  (q.v.)  from  one  another  illustrates  the 
common  case  (see  under  Alcohol  for  the  third  possibility). 

The  composition  of  the  mixture  having  the  minimum  vapor  tension 
varies  with  the  external  pressure,  and  so  does  the  boiling-point.  At  300 
mm.  the  constant  boiling  liquid  contains  21.8  per  cent  of  hydrogen  chlo- 
ride and  boils  at  84° ;  at  1520  mm.  it  contains  19.1  per  cent  of  the  gas. 

The  common  belief  that  hydrochloric  acid  of  constant  boiling-point  is  a  definite 
compound  is  without  foundation.  Compounds  do  not  vary  in  composition  with 
changes  in  pressure  in  this  manner.  Aqueous  solutions  of  hydrogen  iodide,  hydro- 
gen bromide,  and  nitric  acid  behave  in  the  same  way.  But  solutions  of  oxygen,  of 
ammonia,  and  of  many  liquids  (e.g.  methyl  alcohol)  in  water  belong  to  the  second 
of  the  two  classes  mentioned  above,  and  the  more  volatile  component  often  leaves 
the  water  entirely  before  much  of  the  latter  has  evaporated. 

Chemical  Properties.  —  This  compound  is  extremely  stable,  as 
we  might  expect  from  the  vigor  with  which  the  elements  of  which  it  is 
composed  combine.  On  being  heated  to  a  temperature  of  1800°  it  be- 
gins, however,  to  dissociate  into  its  constituents. 

In  the  chemical  point  of  view,  it  is  on  the  whole  rather  an  indif- 
ferent substance.  When  water  is  saturated  with  the  gas  at  —  22°  a 
hydrate  (HC1,  2HjO)  crystallizes  out.  This  decomposes  into  the  same 
constituents  when  allowed  to  warm  up  again  to  —  18°.  Hydrogen 
chloride  (the  gas)  has  no  action  upon  any  of  the  non-metals,  such  as 
phosphorus,  carbon,  sulphur,  etc.  Many  of  the  metals,  however,  par- 
ticularly the  more  active  ones,  such  as  potassium,  sodium,  and  mag- 
nesium, decompose  it.  Hydrogen  is  set  free,  and  the  chloride  of 
the  metal  is  formed  (K  +  HC1  ->  KC1  +  H).  Hydrogen  chloride 
unites  directly  with  ammonia  gas  to  form  solid  ammonium  chloride 
(HC1  4-  NHg  —>  NH4C1).  The  liquefied  gas  has  the  same  properties. 

Composition.  —  The  proportion  of  hydrogen  to  chlorine  by  weight 
in  this  compound  is  1  :  35.18.  Taking  the  atomic  weight  of  hydrogen 
1.008,  so  as  to  harmonize  that  of  chlorine  with  0  =  16,  the  ratio 
becomes  1.008 : 35.45. 


184 


INORGANIC   CHEMISTRY 


The  proportion  by  volume  in  which,  the  constituents  unite,  and 
the  relation  of  this  to  the  volume  of  the  resulting  hydrogen  chloride, 
may  easily  be  shown  in  several  ways.  The  decomposition  of  the  solu- 
tion of  hydrogen  chloride  in  water  by  means  of  the  electric  current 
proves  that  the  gases  are  liberated  in  equal  volumes. 

The  apparatus  in  Fig.  31  (p.  94)  cannot  be  used  to  show  this,  because,  under 
the  increasing  pressure  due  to  the  displacement  of  the  liquid  into  the  higher  bulb, 
the  chlorine  becomes  more  and  more  soluble,  and  its  volume  therefore  falls  pro- 
gressively more  and  more  below  what  it  should  be. 

A   special  form  of  apparatus  (Fig.  67)  was   devised   by  Lothar 
Meyer  to  demonstrate  the  volumetric  proportion.     The  central  part  is 

the  same  as  in  Fig.  31, 
but  the  gases  go  to 
right  and  left,  and  dis- 
place the  liquid  in  two 
inverted  tubes.  The 
equal  rate  at  which 
this  takes  place  on 
both  sides  proves  that 
the  gases  are  generated 
in  equal  volumes. 

In  order  to  ascer- 
tain the  relation  be- 
tween the  volumes  of 
the  constituents  and 

FlQ  67  that  of  the  product,  we 

may  unite  the  gases 
and  find  out  whether 

any  change  in  volume  occurs.  A  tube  with  thick  walls  (Fig.  68) 
is  filled  with  the  mixed  gases  obtained  by  electrolysis.  By  dipping 
one  end  of  the  tube  under  mercury  and  opening  the  lower  stop- 
cock, it  is  seen  that  no  gas  leaves  and  no  mercury  enters.  After 
the  mixture  has  been  exploded,  by  the  light  from  burning  magnesium, 
the  same  test  is  repeated  with  the  same  result.  The  pressure  has 
therefore  remained  equal  to  that  of  the  atmosphere.  Hence  there  has 
been  no  change  in  volume  as  the  result  of  the  union.  It  appears 
therefore,  that : 

1  vol.  hydrogen  +  1  vol.  chlorine  — >  2  vols.  hydrogen  chloride, 
a  result  in  harmony  with  Gay-Lussac's  law  (p,  125). 


CHLORINE   AND  HYDROGEN   CHLORIDE  185 

Another  way  of  demonstrating  the  equality  in  the  volumes  of  the  hydrogen 
and  chlorine  is  to  fill  a  wide  tube,  closed  at  each  end  by  a  stopcock,  with  the  mixed 
gases  arising  from  electrolysis  of  hydrochloric  acid.  When  the  air  has  been 
entirely  displaced,  the  stopcocks  are  closed.  The  gases  which  the  tube  then  con- 
tains are  present  very  nearly  in  the  proportions  in  which  they  are  liberated  from 
the  decomposition  of  the  substance.  By  introducing  a  small  amount  of  potassium 
iodide  solution,  the  chlorine  is  removed.  It  forms  potassium  chloride;  which 
remains  dissolved  in  the  water,  and  free  iodine,  which  dissolves  in  the  excess  of 
potassium  iodide  solution  (Kl  +  Cl  — »KC1  +  I). '  Neither  product  is  gaseous 
under  the  circumstances,  so  that  the  volume  of  the  mixed  gases  dimin- 
ishes by  the  amount  of  chlorine  removed.  If  the  stopcock  is  now 
opened  under  water,  the  latter  enters  and  fills  half  the  length  of  the 
tube.  The  remaining  gas  is  easily  shown  to  be  hydrogen. 

To  show  that  the  volume  of  the  hydrogen  chloride  is  twice  that  of 
either  constituent  in  the  free  condition,  an  alternative  method  is  like- 
wise available.  We  may  completely  fill  a  long  test-tube  with  hydrogen 
chloride,  introduce  into  it  quickly  some  sodium  dissolved  in  mercury, 
and,  after  agitation,  open  the  closed  tube  under  mercury.  The  sodium 
gives  sodium  chloride  and  free  hydrogen,  and  it  is  found  that  the 
mercury  enters  so  as  to  fill  one-half  of  the  tube.  Since  from  this 
experiment  we  learn  that  the  hydrogen  occupies  half  the  space  of  the 
hydrogen  chloride,  and  from  the  previous  experiment  we  know  that  the 
volume  of  hydrogen  is  equal  to  that  of  the  chlorine,  we  conclude  that 
two  volumes  of  mixed  hydrogen  and  chlorine  would  give  two  volumes 
of  hydrogen  chloride. 

Chlorides.  —  The  chlorides  are  described  individually 
under  the  other  element  which  each  contains.  For  the  pres- 
ent we  simply  add  to  the"  statement  on  p.  179  that  the 
majority  of  the  chlorides  which  do  not  interact  with  water 
(p.  181),  that  is  to  say  of  the  chlorides  of  the  metals,  are 
easily  soluble  in  water.  The  only  exceptions  are  silver 
chloride  (AgCl),  mercurous  chloride  (calomel,  HgCl),  cuprous 
chloride  (CuCl),  aurous  chloride  (one  of  the  chlorides  of  gold, 
AuCl),  thallous  chloride  (T1C1),  and  ordinary  lead  chloride 
(PbCl2).  The  last  of  these  is  on  the  border  line  as  regards 
solubility.  An  appreciable  amount  dissolves  in  cold  water,  and  a  con- 
siderable amount  in  boiling  water. 

Chemical  Properties  of  Hydrochloric  Acid.  —  The  solution  of 
hydrogen  chloride  in  water  is  an  entirely  different  substance  in  its 
chemical  behavior  from  hydrogen  chloride.  It  is  strongly  acid,  turn- 
ing litmus  red.  The  gas  and  liquefied  gas  have  no  such  property. 
The  solution  conducts  electricity,  as  we  have  seen,  very  well,  and  is 
decomposed  in  the  process  (cf.  p.  194). 


186  INORGANIC   CHEMISTRY 

Many  metals,  when  introduced  into  hydrochloric  acid,  displace  the 
hydrogen  (p.  95),  and  form  the  chloride  of  the  metal.  In  the  case  of 
zinc  the  action  was  represented  by  the  equation  : 

Zn  +  2HC1  -»  ZnCl2  +  2H. 

The  liquefied  gas  has  no  action  upon  zinc,  and  even  its  solution  in 
many  solvents  shows  little  activity.  The  solution  in  alcohol  behaves 
like  that  in  water.  But  the  solutions  in  toluene,  benzene,  and  other 
compounds  of  carbon  and  hydrogen,  in  many  of  which  the  gas  is 
freely  soluble,  are  hardly  affected  by  the  presence  of  zinc  and  other 
metals.  These,  and  many  other  facts  which  we  shall  notice  later 
(see  Dissociation  in  solution),  show  that  the  condition  of  this  sub- 
stance in  aqueous  solution  is  peculiar. 

The  aqueous  solution  of  hydrogen  chloride  interacts  rapidly  with 
most  oxides  and  hydroxides  of  metals,  as,  for  example,  those  of  zinc : 

ZnO  +  2HC1  ->  ZnCL,  +  H2O, 
Zn(OH)2  +  2HC1  ~*  ZnCl2  +  2H2O. 

Here  no  free  hydrogen  is  obtained,  since  the  oxygen  in  the  oxide,  and 
the  hydroxyl  in  the  hydroxide,  unite  with  it  to  form  water.  In  each 
case,  however,  the  chloride  of  the  metal  is  obtained.  It  may  be  noted 
in  passing  that  all  acids  behave  in  a  similar  manner  towards  oxides 
and  hydroxides,  giving  water  and  a  compound  corresponding  to  the 
chloride  (cf.  p.  171).  Dilute  sulphuric  acid,  for  example,  gives  sulphates. 

In  the  two  preceding  paragraphs,  three  kinds  of  actions,  each  con- 
stituting a  different  way  of  obtaining  chlorides,  have  been  mentioned 
incidentally.  There  are  two  others  which  we  have  already  en- 
countered. The  simplest  is  the  direct  union  of  the  element  with 
chlorine  (Zn  +  201  — >  ZnCl2).  The  other  method  is  illustrated  in  the 
case  of  the  precipitation  of  silver  chloride  (p.  13).  Here  the  forma- 
tion of  the  chloride  occurred  by  exchange  of  another  radical  for 
chlorine  (AgNO8  +  NaCl  -*  AgClJ  +  NaN08).  The  insoluble  chlo- 
rides (p.  185)  can  be  made  conveniently  by  this  plan.  The  formation 
of  the  precipitates,  for  example  that  of  silver  chloride,  is  used  as  a 
test  for  the  presence  of  a  soluble  chloride  in  the  solution. 

The  solution  of  hydrogen  chloride  in  water  sold  in  commerce  is 
known  by  the  name  of  muriatic  acid  (Lat.  muria,  brine).  It  is  a 
yellow  liquid  which  contains  a  number  of  impurities.  The  most  com- 
mon are  ferric  chloride,  which  is  responsible  for  part  of  the  yellow 
tint,  some  yellow  organic  coloring  material,  arsenious  chloride,  and 


CHLORINE  AND   HYDROGEN  CHLORIDE  IbV 

free  chlorine.  The  acid  frequently  gives  a  residue  when  evaporated, 
and  this  must  of  course  represent  some  impurity.  It  is  sometimes 
adulterated  with  calcium  chloride,  since  the  price  obtained  depends 
upon  the  specific  gravity  of  the  solution,  and  this  may  be  raised  by 
dissolving  calcium  chloride  in  it. 

Classification  of  Chemical  Interactions  and  Exercises 
Thereon.  —  So  far  we  have  defined  ten  more  or  less  distinct  kinds 
of  chemical  change :  Combination  (p.  14),  decomposition  (p.  15), 
dissociation  (p.  121),  displacement  (p.  99),  substitution  (p.  176), 
double  decomposition  (p.  99),  hydrolysis  (p.  181),  oxidation  (pp. 
72,  110,  172),  reduction  (pp.  72,  110),  and  electrolysis  (p.  19).  In  one 
or  two  of  these  classes  all  the  actions  are  reversible,  in  others  some  are 
reversible  and  some  are  not.  Illustrations  of  every  one  of  these  will  be 
found  in  the  present  chapter.  The  classes  are  not  mutually  exclusive. 
Some  actions  belong  to  one  class  or  another  according  to  our  point  of 
view  at  the  moment.  The  ability  readily  to  classify  each  phe- 
nomenon, as  it  comes  up,  requires  precisely  that  grasp  of  the  frame- 
work of  the  science  which  the  reader  must  seek  speedily  to  attain. 
For  example,  let  him  classify  the  following  actions :  1.  Action  of 
heat  on  chloroplatinic  acid ;  2.  of  potassium  on  water ;  3.  of  heat  on 
potassium  chlorate ;  4.  of  chlorine  on  metals ;  5.  of  chlorine  on  tur- 
pentine ;  6.  of  chlorine  on  potassium  iodide ;  7.  of  chlorine  on 
methane ;  8.  of  carbon  monoxide  and  chlorine ;  9.  of  sunlight  on 
hypochlorous  acid ;  10.  of  sulphuric  acid  on  salt ;  11.  of  zinc  oxide 
and  hydrochloric  acid ;  12.  of  zinc  on  hydrochloric  acid. 

13.  Expand  the  explanation  of  the  tendency  of  hydrogen  chloride 
to  fume  in  moist  air  (p.  182). 

14.  Explain  the  interaction  of  steam  and  iron  (p.  110)  on  mechanical 
principles  similar  to   those  used    in   describing  how  hydrogen   chlo- 
ride is  formed  from  salt  and  sulphuric  acid  (p.  179). 

15.  What  is  the  maximum  that  the  temperature  of  the  action  (p.  172) 
of  hydrochloric  acid  on  potassium  permanganate  may  attain  ?     Why 
is  this  interaction  so  much  more  vigorous  than  that  where  manganese 
dioxide  is  used  (p.  173)  ? 

16.  In  view  of  the  explanations  given,  can  you  define  the  general 
nature  of  the  "  other  substances  "  (p.  172)  which  may  be  used  to  oxi- 
dize hydrochloric  acid  ? 


188  INORGANIC  CHEMISTRY 

SUMMARY  OF  PRINCIPLES. 

It  may  be  useful  at  this  point  partially  to  summarize  the  principles  (general 
facts)  of  chemistry  so  far  as  they  have  been  developed  in  the  preceding  chapters. 
These  principles  are  given  under  fourteen  heads  below.  They  are  stated  as  far 
as  possible  strictly  in  terms  of  facts,  since  hypotheses  are  not  integral  parts  of 
chemistry,  but  are  scaffolding  temporarily  employed  to  facilitate  the  erection  of 
the  structure  of  the  science.  In  a  later  chapter  (Chap,  xv),  some  other  impor- 
tant principles  will  be  summarized  in  like  manner.  To  secure  more  strictly  logical 
arrangement  than  has  seemed  advisable  in  the  text,  two  conceptions  which  have 
already  been  dealt  with  are  held  over  to  the  second  half  of  the  summary,  namely, 
17  (valence)  and  21  (chemical  relations  of  elements).  The  reader  should  give  care- 
ful thought  to  the  various  points,  many  of  which,  in  a  backward  view,  will  be  found 
to  have  become  susceptible  of  improved  statement.  We  begin  the  series  with  the 
most  fundamental  fact  of  all,  —  the  one  without  which  no  chemical  work  would  be 
possible  : 

1.  Each  substance  has  its  own  set  of  specific  physical  properties.     By  means 
of  these  it  is  recognized  and,  when  necessary,  separated  from  other  substances  (p.  5). 

2.  Substances  are  either  simple   (elementary),  containing  only  one  kind  of 
matter,  or  compound,  containing  more  than  one  kind  of  matter  (p.  30). 

3.  In  all  chemical  phenomena  (excepting  u  internal  rearrangements  "),  changes 
in  the  material  composition  of  bodies  occur  (p.  16). 

4.  In  chemical  phenomena  there  is  no  change  in  the  total  mass  of  the  system 
(P-  17). 

6.   Each  substance  has  a  definite  material  composition  by  weight  (p.  41). 

6.  The  proportions  by  weight  in  which  all  chemical  combinations  take  place 
can  be  expressed  in  terms  of  small  integral  multiples  of  fixed  numbers,  which  may 
be  called  combining  weights,  one  for  each  element.     That  weight  of  each  element 
which  combines  with  8  parts  of  oxygen  is  called  the  equivalent  weight  and  has  the 
properties  of  a  combining  weight  (pp.  48,  50). 

7.  The  proportions  by  volume  in  which  all  chemical    interactions  involving 
substances  in  the  gaseous  condition  take  place  are  expressible  by  small  integers 
(p.  125). 

8.  From  6  and  7  it  follows  that  the  equivalent  weights  of  all  substances  occupy, 
in  the  gaseous  condition  and  at  the  same  temperature  and  pressure,  volumes  which 
are  either  equal  or  stand  to  one  another  in  the  ratio  of  small  integers  (p.  126). 

9.  In  every  chemical  phenomenon  a  transformation  of  energy  occurs.     This 
results  in  a  redistribution  of  the  chemical  energy  in  the  substances  concerned,  and 
also  in  an  increase  or  a  decrease  in  the  total  chemical  energy  in  the  system  (p.  26). 

DEFINITIONS  :  The  word  substance  is  applied  to  a  compound  of  one  or  more 
kinds  of  elementary  matter  with  a  certain  proportion  of  chemical  energy  (p.  82). 
The  word  element  is  applied  to  a  variety  of  simple  matter  which  exists  only  in  com- 
bination with  energy,  and  often  with  other  kinds  of  elementary  matter  as 
well  (p.  31). 

10.  In  chemical  phenomena  there  is  no  actual  loss  or  gain,  but  only  transfor- 
mation of  energy  (p.  23). 

11.  Interactions  which  proceed  spontaneously  are  in  general  those  in  which 
the  free  energy  is  transformed   into  some  other  variety  or  varieties   of  energy 
(p.  27). 


CHLORINE   AND   HYDROGEN   CHLORIDE  189 

12.  Each  substance  has  its  own  set  of  chemical  properties,  such  as  : 

(a)  Affinity:  the  given  substance  can  or  can  not  interact  with  such  and  such 
elementary  and  compound  substances. 

(6)  Relative  activity  of  the  systems  in  (a).  This  is  measured  quantitatively 
by :  (a)  Relative  speed  under  like  conditions  (see  18,  14,  18);  (/3)  Relative  heat 
developed,  when  actions  compared  can  be  carried  out  so  that  all  conditions  are 
alike ;  (7)  Relative  E.M.F.  of  cell  when  the  action  is  so  arranged  as  to  give 
electricity  (pp.  28,  76,  111). 

13.  The  speed  of  every  interaction  is  increased  by  raising  the  temperature  (p.  72). 

14.  The  speed  of    interactions  is  increased  or  decreased  by  catalytic  agents, 
each  of  which  is  individual  in  the  kind  and  amount  of  its  effect  (p,  74), 


CHAPTER   XII 
MOLECULAR  WEIGHTS   AND   ATOMIC  WEIGHTS 

AVOGADRO'S  hypothesis  (p.  131)  has  proved  to  be  by  far  the 
most  suggestive  and  fruitful  of  all  the  conceptions  developed  from 
the  kinetic-molecular  hypothesis.  We  are  now  in  a  position  to 
discuss  several  of  its  most  important  applications.  To  speak  in 
terms  of  the  hypothesis,  these  concern  more  particularly  the  measure- 
ment of  the  relative  weights  of  the  molecules  of  different  gaseous 
substances,  and  the  determination  of  the  most  convenient  magni- 
tudes for  the  chemical  unit  weights  (atomic  weights ;  cf.  p.  50). 

Meaning  of  Avogadro's  Hypothesis.  —  First,  we  must  under- 
stand clearly  what  is  implied  in  the  statement  that :  In  equal  volumes 
of  all  gases,  at  the  same  temperature  and  pressure,  there  are  equal 
numbers  of  molecules.  It  means  that,  for  instance,  at  100°  and  760 
mm.,  in  all  specimens  of  gases  the  average  spacing  of  the  molecules 
is  identical.  This  condition  is  independent  of  the  nature  of  the  gas 
—  for  example,  whether  it  is  a  simple  or  a  compound  substance,  like 
oxygen  and  carbon  dioxide  respectively,  or  a  mixture,  like  air.  It 
means  that  when,  at  some  fixed  temperature,  we  fill  the  same  vessel 
with  a  number  of  different  gases  or  gaseous  mixtures  successively, 
the  number  of  molecules  that  it  will  hold  at  a  pressure,  say,  of  one 
atmosphere  will  always  be  the  same.  If  we  take  care  to  keep  tem- 
perature and  pressure  the  same,  the  equality  in  the  number  of  mole- 
cules that  will  enter  the  jar  will  take  care  of  itself  automatically. 
In  what  follows,  to  avoid  continual  repetition,  it  is  to  be  assumed 
that  temperatures  and  pressures  are  equal  unless  the  contrary  is 
expressly  stated. 

This  statement  would  be  strictly  true  only  in  the  case  of  gases,  if  such  existed, 
which  behaved  in  ideal  accord  with  the  laws  of  Boyle  and  Charles.  Since,  how- 
ever, in  all  gases,  with  the  exception  of  hydrogen,  a  certain  tendency  to  cohesion 
between  molecules  is  distinctly  noticeable  and  its  amount  varies  from  gas  to  gas, 
the  density  with  which  the  molecules  are  packed  is  not  precisely  the  same  in  any 
two  of  them  (p.  132).  Hence,  Avogadro's  hypothesis  is  not  perfectly  realized  in 
any  known  gases.  In  the  case  of  hydrogen,  for  example,  a  divergence  from  the 
behavior  of  an  ideal  gas  exists,  but  is  barely  measurable,  and  in  the  case  of 

190 


MOLECULAR   WEIGHTS 


191 


chlorine,  it  amounts  to  about  1£  per  cent,  and  is  quite  conspicuous.  This  slight 
irregularity  in  the  packing  of  the  molecules,  however,  does  not  interfere  with 
the  application  of  this  hypothesis  in  chemistry. 

MOLECULAR   WEIGHTS. 

The  Relative  Weights  of  the  Molecules.  —  According  to  Avo- 
gadro's  hypothesis,  vessels  of  equal  size  filled  with  different  gases 
contain  equal  numbers  of  gaseous  molecules.  Now  equal  volumes  of 
different  gases  differ  very  markedly  in  weight,  or,  in  other  words, 
the  densities  of  various  known  gases  cover  a  wide  range  of  values. 
Thus,  hydrogen  is  the  lightest  of  all,  chlorine  is  more  than  thirty- 
five  times,  mercuric  chloride  (corrosive  sublimate)  vapor  over  one 
hundred  and  thirty-four  times  as  heavy.  Since  these  different 
weights  of  equal  volumes  represent  the  weights  of  equal  numbers  of 
molecules,  the  difference  must  be  due  to  the  differing  weights  of  the 
molecules  themselves.  The  densities  of  gases,  therefore,  may  be  taken 
as  measures  of  the  relative  weights  of  their  individual  molecules.  The 
extreme  significance  of  this  inference  in  chemistry  will  appear  as  we 
elaborate  upon  it. 

The  various  scales  on  which  the  densities  of  gases  may  be  calcu- 
lated, such  as  the  weights  of  one  liter  of  each  gas,  or  the  weights  of 
volumes  equal  to  that  of  one  gram  of  air,  are  illustrated  in  the  first 
two  columns  of  the  following  table  : 


Weight  of 
One  Liter,  0° 
and  760  mm. 

Density, 
Air  =  1. 

Molecular 
Weight, 
Ox.  =  32. 

Hydrogen  

0.090 

0.0696 

2.016 

Oxygen  
Chlorine  
Hydrogen  chloride  .  .  . 

1.429 
3.166 
1  628 

1.105 
2.449 
1  259 

32.00 

70.90 
36  458 

Carbon  dioxide  
Water  .  .  

1.966 
0  8045 

1.520 
0  622 

44.00 
18  016 

Mercury  
Mercuric  chloride  
Air  .  

8.932 
12.097 
1.293 

6.908 
9.354 
1.00 

200.0 
270.90 
28  955 

The  values  for  water  (b.-p.  100°),  mercury  (b.-p.  357°),  and  mer- 
curic chloride  (b.-p.  305°)  are  measured  at  high  temperatures  and 
reduced  by  rule  (p.  91)  to  0°  and  760  mm.  All  the  numbers  in  the 
first  two  columns,  as  they  stand,  are  purely  physical  in  derivation. 
Those  in  the  second  column  are  obtained  from  those  in  the  first  by 
using  the  proportion : 


192 


INORGANIC   CHEMISTRY 


1.293  (wt.  1 1.  air)  :  1.00  (air  =  1)  : :  wt.  of  1  1.  any  gas  :  x  (dens,  of 
that  gas). 

The  last  column  will  be  explained  presently.  Since  the  numbers  in 
the  first  column  apply  to  equal  volumes  (1  1.),  and  those  in  the 
second  stand  in  constant  ratios  to  them,  the  weights  in  the  second 
column  represent  equal  volumes  also.  In  the  second,  the  volume 
is  ^jf-y  1.  The  values  in  either  one  of  the  columns  represent  the 
relative  weights  of  the  molecules  of  the  various  substances  (see 
Exercise  1  in  this  chapter). 

In  order  to  avoid  the  creation  of  unnecessary  confusion  in  the  mind  of  the 
beginner,  the  weights  of  one  liter  of  gas  in  the  above  table  are,  with  the  exception 
of  that  of  oxygen,  all  ideal  numbers.  They  are  calculated  back  from  the  correct 
molecular  weights  made  up  from  the  atomic  weights.  This  enables  us  to  show  the 
process  by  which  the  molecular  weights  are  derived  from  the  weights  of  one  liter 
without  the  exhibition  of  arithmetical  discrepancies  which  might  obscure  the 
principle  being  explained.  The  weights  of  one  liter  of  the  various  gases,  as  we 
have  given  them,  are  based  on  the  assumption  that  the  molecules  are  always 
packed  uniformly  in  accordance  with  Avogadro's  hypothesis.  The  actual  values 
are  in  most  cases  somewhat  different  from  these,  and  we  attribute  the  divergencies 
to  the  varying  degrees  of  cohesion  between  the  molecules  of  different  substances. 
Even  the  weight  of  one  liter  of  the  same  gas,  after  reduction  to  0°  and  760  mm., 
is  found  to  vary  with  the  temperature  and  pressure  at  which  it  was  examined. 
This  is  but  natural,  since  changes  in  these  conditions  alter  the  effects  of  cohesion. 
The  following  table  gives  the  actual  weights  of  one  liter  of  the  same  gases,  with  a 
few  additional  ones,  and  a  comparison  will  show  the  extent  of  the  divergencies. 
The  most  interesting  case  perhaps  is  that  of  oxygen  and  hydrogen.  The  chemical 
combining  weights  of  these  substances  are  in  the  ratio  of  15.88  : 1.00,  while  a 
slight  excess  of  cohesion  in  oxygen  gives  the  ratio  of  their  densities  the  value 
15.90  :  1.00.  These  numbers  are  in  both  cases  based  upon  Morley's  results. 


Weight  of 
One  Liter,  0° 
and  760  mm. 

Density, 
Air  =  l. 

Observed  Mole- 
cular Weight, 
Ox.  =  32. 

Adjusted 
Molecular 
Weight. 

Hydrogen  ...... 

0  08987 

0  0695 

2  012 

2  016 

Oxvsren 

1.429 

1  105 

32  OO 

3200 

Nitrogen     

1.2507 

0  967 

28.07 

28.08 

Chlorine     

3.220 

2.490 

72.01 

70.90 

Hydrogen  chloride  .     .     . 
Carbon  dioxide    .... 
Hydrogen  sulphide  .     .     . 
Ammonia  .               ... 

1.6398 
1.9768 
1.537 
0.7708 

1.269 
1.529 
1.189 
0.597 

36.72 
44.27 
34.43 
17.26 

36.458 
44.00 
34.076 
17.064 

Sulphur  dioxide  .... 
Water                  .... 

2.9266 
08046 

2.264 
0  622 

65.54 
18  016 

64.06 
18  016 

Mercury     

8.87 

6.86 

198.4 

200.0 

Air    

1.293 

1.00 

28.955 

[Mixture] 

MOLECULAR   WEIGHTS  193 

Molecular  Weights.  —  The  foregoing  section  shows  that,  provided 
a  substance  is  a  gas  or  can  be  volatilized,  the  relative  weight  of  its 
molecules,  compared  with  those  of  other  volatile  substances,  can  be 
ascertained.  To  save  words,  the  relative  weight  of  the  molecules  of  a 
substance  is  called  the  molecular  -weight  of  the  substance.  Since  the 
absolute  weights  of  molecules  cannot  be  determined,  the  next  question 
which  arises  is  as  to  the  choice  of  an  appropriate  unit,  and  therefore 
an  appropriate  scale  for  these  relative  weights,  or  molecular  weights. 
Now  the  numbers  already  given,  in  the  first  two  columns  of  the  table, 
are  purely  physical  data  and,  as  they  stand,  lack  direct  relation  to 
chemical  facts.  A  set  of  chemical  numbers  is  required  for  chemical 
purposes.  We,  therefore,  proceed  next  to  show  how  the  required 
relation  between  the  relative  weights  of  molecules  and  chemical  facts 
can  be  established. 

Chemistry  deals  with  chemical  combination,  and  most  substances 
are  compounds.  If  we  fix  our  attention,  then,  first,  on  compound 
substances  and  their  molecules,  we  perceive  at  once  that  the  mole- 
cules of  a  compound  substance  must  contain  two  or  more  elements 
in  definite  proportions  by  weight.  We  may  therefore  extend  the 
molecular  hypothesis  by  supposing  that  a  molecule  is  composed  of 
smaller  parts,  which  we  call  atoms.  The  atoms  will  be  elementary. 
Thus  the  molecule  of  the  compound  will  contain  one  or  more  atoms 
of  each  of  the  component  elements.  This  conception,  the  starting- 
point  of  the  atomic  hypothesis,  is  elaborated  in  the  next  chapter. 
Its  introduction  in  the  present  connection,  however,  at  once  suggests 
two  ideas.  In  the  first  place,  since  we  can  now  determine  the  relative 
weights  of  molecules,  we  should  also  somehow  be  able,  with  the  help 
of  the  combining  proportions,  to  determine  the  relative  weights  of 
the  atoms  of  the  Elements.  If  these  relative  weights  of  atoms  are 
properly  determined,  then  the  weights  of  the  atoms,  when  added 
together,  should  give  the  weight  of  the  molecule.  Thus, — and  this 
is  the  second  idea,  and  the  one  of  most  immediate  use  to  us,  —  the  scale 
for  relative  weights  of  molecules  must  be  chosen  with  reference  to 
the  scale  for  relative  weights  of  the  atoms,  so  that  the  former  weights 
may  always  include  the  latter.  That  is  to  say,  the  molecular  weights 
must  be  based  upon  the  combining  weights.  This  means'  that  the 
scale  must  be  such  that  no  molecule  of  a  compound  of  hydrogen  shall 
receive  a  value  so  small  that  the  proportion  of  hydrogen  in  it  is  less 
than  1.008.  Furthermore,  since  the  combining  weight  of  oxygen  is 
the  standard  for  combining  proportions,  it  is  desirable  to  use  this 


194  INORGANIC   CHEMISTRY 

substance  as  basis  of  the  scale  of  molecular  weights.  Now,  as  we 
shall  find  (see  p.  197),  it  turns  out  that,  to  avoid  obtaining  propor- 
tions of  hydrogen  less  than  unity,  we  are  compelled  to  take  the  scale 
32  for  the  molecular  weight  of  oxygen.  Our  chemical  scale  of  densi- 
ties is  therefore  calculated  to  the  scale,  density  of  oxygen  =  32.  This, 
then,  is  the  answer  to  the  problem  with  which  the  section  opened. 
The  third  column  of  the  table  (p.  191)  shows  the  results  of  recalculat- 
ing the  densities  to  this  chemical  scale.  The  proportion  used  is  : 

Density  of  Ox.  :  Density  of  Substance  :  :  32 :  x. 

Thus,  if  we  take  the  densities  from  the  first  column,  the  value  for 
water  is  found  by  the  proportion,  1.429  :  0.8045  : :  32  :  x  (=18.016). 
That  is,  we  multiply  the  weight  of  a  liter  of  the  gas  by  32/1.429  to 
get  the  molecular  weight. 

Since  the  gram  is  the  unit  of  weight,  32  g.  of  oxygen,  or  18.016  g. 
of  water  is  called  the  gram-molecular  weight  of  the  substance.  It 
will  be  noted  that  32  g.  is  not  the  weight  of  a  molecule  of  oxygen. 
It  is  the  weight  of  a  very  large,  and  not  exactly  known  number  of 
molecules  of  oxygen.  But,  whatever  that  number  of  molecules  of 
oxygen  is,  18.016  g.  of  water  contains  the  same  number  of  molecules, 
and  the  other  weights  in  the  same  column  are  weights  of  numbers  of 
molecules  equal  to  these.  The  term  gram-molecular  weight  being 
somewhat  ponderous,  we  abbreviate  it  to  molar  -weight,  and  still  fur- 
ther to  mole.  Thus,  a  mole  of  chlorine  is  70.9  g.  of  the  element,  and 
a  mole  of  hydrogen  chloride  is  36.458  g.  of  the  compound. 

When  the  above  method  of  calculating  the  molecular  weight  from  the  weight 
of  one  liter  of  a  substance  is  applied  to  the  actual  experimental  values,  the  result- 
ing molecular  weights  necessarily  diverge  somewhat  from  the  ideal  ones  which  we 
have  given.  Thus,  since  a  liter  of  hydrogen  chloride  actualjy  weighs  1.6398,  this 
number  when  multiplied  by  32  and  divided  by  1.429  gives  the  value  36.72.  This 
value  of  the  molecular  weight,  however,  does  not  contain  the  atomic  weight  of 
chlorine,  as  found  by  measuring  combining  proportions,  an  even  number  of  times. 
Consequently,  the  molecular  weights  actually  determined  by  experiment  have 
always  to  be  adjusted  by  a  slight  numerical  change  to  the  nearest  value  which 
contains  integral  multiples  of  the  atomic  weights  of  the  constituents.  In  the  table 
on  page  192  the  observed  and  the  ad  justed  molecular  weights  have  both  been  given. 
The  mole  is  always  the  adjusted  molecular  weight  and  not  the  observed  one. 

This  process  of  adjustment  does  not  really  involve  any  uncertainty,  because  it 
never  requires  changes  large  enough  to  introduce  actual  confusion.  The  adjust- 
ment of  the  observed  molecular  weight  to  correspond  with  the  measured  combin- 
ing weights  is  made  in  preference  to  the  converse  operation  because  the  latter 
values  are  always  susceptible  of  exact  determination  while  the  former  are  not. 
We  cannot  expect  the  molecular  weights  to  be  accurate,  because,  in  the  first  place, 


MOLECULAR   WEIGHTS 


196 


the  measurement  of  the  weight  and  volume  of  a  gas  or  vapor  is  always  difficult, 
and  often  involves  an  error  of  one-half  per  cent,  and,  in  the  second  place,  the 
way  in  which  the  behavior  of  gases  departs  from  that  of  a  perfect  gas  causes 
the  results  to  vary  according  to  the  temperature  and  pressure  chosen  for  the 
experiment. 

The  Gram- Molecular  (Molar)  Volume.  —  The  weights  in  the 
last  column  of  the  table  (p.  191)  must  represent  equal  volumes  of  the 
different  gases.  This  follows  from  the  fact  that  they  are  derived 
from  the  values  in  the  first  column  by  multiplying  by  a  constant  ratio 
(32/1.429),  and  the  volume  in  the  first  column  is  always  1  liter. 
The  actual  dimension  of  this  volume  is  evidently  32/1.429  liters, 
which  is  almost  exactly  22.39,  or  in  round  numbers  22.4  liters.  This 
volume  at  0°  and  760  mm.  holds  32  g.  of  oxygen,  70.9  g.  of  chlorine, 

44.00  g.  of  carbon  dioxide,  or,  in  fact,  the  molar  weight  of  any  gaseous 
substance.     It  is  called,  therefore,  the  gram-molecular  volume  (G.M.V.) 
or  the  molar  volume.    It  may  be  defined  as  that  volume  which  contains 
one  mole  (gram-molecular  weight)  of  any  gas  at  0°  and  760  mm.     At 
other   temperatures   and   pressures  the    G.M.V.  has    correspondingly 
different  values. 

The  G.M.V.  gives  us  a  concrete  conception  of  a  molar  weight. 
This  volume  is  represented  by  a  cube  (Fig.  69)  28.19  cm.  (or  about 

11.1  inches)  high.     Like  any  other  vol- 
ume, it  holds  identical  numbers  of  mole- 
cules of  different  gases.*     Its   capacity 
at  0°  and   760  mm.   is  the  number  of 
molecules  in  32  g.  of  oxygen.     Hence, 
in  terms  of  the  hypothesis,  the  weight 
of  any  gas  which  fills  it  bears  to  32  g.    _ 
the  same  ratio  as  the  weight  of  a  mole- 
cule  of  that   gas    to   the   weight   of  a 
molecule   of   oxygen.     We   may,  there- 
fore, state  the  method  of  finding  the  molar  (gram-molecular)  weight 
of  a  substance  thus :  Weigh  a  known  volume  of  the  substance,  at  any 
temperature  and  pressure  at  which  it  is  gaseous,  reduce  this  volume  by 
rule  to  0°  and  760  mm.,  and  calculate  by  proportion  the  weight  of  22.4 
liters  (see  Exercises  1,  2,  3,  5). 

*  A  common  question  is  :  Do  not  molecules  of  different  substances  differ  in 
size,  and  will  not  the  numbers  required  to  fill  the  G.M.V.  therefore  be  different  ? 
The  answer  is  that  the  molecules  are  all  so  small  compared  with  the  spaces 
between  them  (at  760  mm.)  that  the  distances  from  surface  to  surface  are  practi- 
cally the  same  as  from  center  to  center.  A  G.M.V.  of  oxygen,  when  liquefied, 


G.M.V. 

22.4  LITERS 


196  INORGANIC  CHEMISTRY 

That  quantity  of  each  substance  which  at  0°  and  760  mm.  would  fill 
the  6.M.V.  cube  is  the  unit  quantity  of  the  substance  for  all  theoretical 
purposes  in  chemistry.  It  represents  the  relative  weight  of  the  mole- 
cules of  the  substance.  We  shall  employ  it  at  once  for  the  purpose  of 
determining  the  relative  weights  of  atoms,  or  atomic  weights. 

It  is  evident  that  the  chemical  molecular  weights,  adjusted  so  as  to  include 
whole  numbers  of  combining  weights,  will  not  all  occupy  exactly  equal  volumes. 
They  represent  exactly  equal  numbers  of  molecules,  and  the  slight  differences  in 
the  closeness  with  which  the  molecules  are  packed  (p.  190)  will  cause  the  values  of 
the  molar  volumes  to  differ  from  gas  to  gas.  The  values  of  this  volume  calculated 
from  the  actual  weights  of  one  liter  of  each  gas  are  as  follows:  Hydrogen,  22.40; 
oxygen,  22.39  ;  nitrogen,  22.45  ;  chlorine,  22.01 ;  hydrogen  chloride,  22.23  ;  car- 
bon dioxide,  22.26;  water,  22.39;  mercury,  22.55.  The  average  value  of  this 
volume  in  the  case  of  the  more  nearly  perfect  gases  is  22.4  liters,  and  this  is, 
therefore,  the  number  which  we  have  used  in  our  definition. 

ATOMIC  WEIGHTS. 

Determination  of  the  Atomic   Weight  of  Each  Element.— 

If  the  paragraphs  dealing  with  combining  weights  are  now  re-read 
(pp.  45-51),  it  will  be  found  that  the  foundations  for  a  system  of 
weights  was  worked  out,  but  that  no  basis  for  definitely  fixing  the 
individual  values  was  discovered.  At  the  time,  the  only  information 
we  had  was  obtained  by  analyzing  compounds  and  reasoning  about  the 
results,  and  evidently  something  more  was  needed  for  the  absolute 
determination  of  the  values.  Thus  on  p.  51,  it  is  pointed  out  that  the 
equivalent  weight,  or  any  multiple  of  it  by  an  integer,  will  serve  for 
expressing  the  proportions  used  by  the  element  in  combining  with 
other  elements.  In  the  history  of  chemistry,  the  uncertainty  as  to 
which  were  the  best  members  to  choose  caused  much  controversy 
among  chemists  and,  for  years,  endless  confusion.  It  was  only  when 
the  comparison  of  the  combining  weights  with  the  relative  weights  of 
molecules  was  at  last  rigorously  applied,  in  the  fashion  now  to  be 
shown,  that  perfect  order  in  chemical  weights  and  formulae  was  finally 
achieved. 

To  determine  the  atomic  weights,  the  plan  of  procedure  is  per- 

gives  less  than  32  c.c.  of  liquid  oxygen,  or  less  than  1/700  of  the  volume  as  gas. 
And  there  are  still  spaces  between  the  molecules  of  the  liquid.  It  is  only  when 
gases  are  so  severely  compressed  that  the  nearness  of  the  molecules  to  one  another 
approaches  that  found  in  the  liquid  condition  that  the  effects  of  the  bulk  of  the 
molecules  become  conspicuous,  and  a  difference  in  the  behavior  of  different  gases 
is  noticeable.  But  in  the  kind  of  chemical  work  discussed  iu  this  chapter,  pres- 
sures over  one  atmosphere  are  not  used. 


ATOMIC    WEIGHTS  197 

fectly  simple.  lu  the  preceding  section  we  settled  upon  the  chemical 
unit  quantity  of  each  substance.  This  is  the  quantity  which,  in  the 
gaseous  condition,  would  fill  the  G.M.V.  (22.4  liters)  at  0°  and  760 
mm.  Now,  we  seek  the  chemical  unit  quantities  of  the  elements  com- 
bined in  each  substance.  Evidently  the  logical  and  consistent  plan 
must  be  to  take  the  amount  of  each  substance  which  fills  the  G.M.V. 
and  find  out  how  much  of  each  element  present  is  contained  in  this 
unit  amount  of  the  substance.  In  other  words,  to  put  the  matter  con- 
cretely, we  imagine  ourselves  filling  the  cube  (Fig.  69)  with  one  com- 
pound after  another,  and  in  each  case  determining  by  analysis  the 
weight  of  each  constituent  element  present  in  a  cube-full  of  the  sub- 
stance. To  carry  out  this  plan,  two  experimental  operations  are 
necessary  with  each  substance  : 

First  we  determine  the  density,  and  this  gives  us  the  gram-molec- 
ular weight,  i.e.  the  amount  filling  the  cube.  This  shows  the  rela- 
tive weight  of  a  molecule  of  the  substance,  as  compared  with  that  of 
one  molecule  of  oxygen. 

Then  we  analyze  the  substance,  and  this  gives  us  the  quantity  of 
each  constituent  in  the  total  gram-molecular  weight,  i.e.  in  the  ma- 
terial filling  the  cube.  This,  in  turn,  shows  the  weight  of  the  quan- 
tity of  each  element  present  in  each  molecule,  relative  to  the  weight 
of  a  molecule  of  oxygen. 

For  example,  the  cube  holds  36.458  g.  of  hydrogen  chloride,  and 
this  amount,  when  decomposed,  yields  1.008  g.*  of  hydrogen  and 
35.45  g.  of  chlorine. 

Finally,  to  determine  the  best  combining  weight  for  a  given  ele- 
ment, we  repeat  the  two  foregoing  operations  with  as  many  different 
compounds  of  the  element  as  possible,  and  then  we  examine  the 
various  quantities  of  the  element  found  in  the  G.M.V.  of  the  various 
compounds.  From  inspection  of  these  quantities  we  quickly  select 
the  value  of  which  all  are  multiples,  by  unity  or  some  integral  num- 
ber. This  value  for  the  combining  weight  is  the  one  accepted.  In 
terms  of  the  hypothesis,  this  is  the  weight  of  one  atom  of  the  element, 
compared  with  the  weight  of  a  molecule  of  oxygen,  and  molecules 
containing  more  than  this  proportion  contain  two,  three,  or  more  atoms 
of  the  element. 

*  It  will  be  observed  that  if  the  unit  for  molecular  weights  had  been  less  than 
the  number  of  molecules  in  32  g.  of  oxygen,  then  an  equal  number  of  molecules  of 
hydrogen  chloride  would  have  contained  less  than  1.008  g.  of  hydrogen,  and  the 
atomic  weight  of  this  element  would  then  have  been  less  than  unity. 


198 


INORGANIC   CHEMISTRY 


For  example,  if  we  are  seeking  the  atomic  weight  of  chlorine,  we 
set  down  the  result  for  hydrogen  chloride  just  given.  Then  we  take 
another  compound  of  chlorine,  say  phosphorus  oxychloride.  We 
determine  the  weight  of  a  measured  volume  of  its  vapor,  at  a  prop- 
erly chosen  temperature  and  pressure,  and  the  result  gives  us,  by 
calculation,  the  molecular  weight,  viz.  153.35.  That  is,  153.35  g.  of 
the  substance  would  fill  the  cube,  if  it  could  be  kept  as  vapor  at  0° 
and  760  mm.  And  this  amount  of  the  substance  contains  31  g.  of 
the  element  phosphorus,  16  g.  of  the  element  oxygen,  and  106.35  g. 
of  the  element  chlorine.  We  then  continue  the  processes  described, 
using  all  the  volatile  compounds  of  chlorine.  The  involatile  com- 
pounds (like  common  salt)  must  be  set  aside,  for  they  cannot  be 
vaporized,  and  therefore  their  molecular  weights  cannot  be  deter- 
mined. When  we  have  studied  as  many  compounds  as  possible  in 
this  way,  we  find  that  there  are  different  quantities  of  chlorine  in 
our  list,  but  they  are  all  integral  multiples  of  35.45  g.  In  phosphorus 


Substance. 

Molar 
Weight 

Weights  of  Constituents  in 
Molar  Weight. 

Hydrogen. 

Chlo- 
rine. 

Oxygen. 

Phosphorus. 

1 

6 

12 
24 
24 
12 
24 

Mercury. 

Molecular 
Formula. 

Hydrogen  chloride      

36.45 
67.45 
137.35 
153.35 
284 
34 
18 
16 
26 
28 
30 
60 
235.45 
270.9 

1 
3 

2 

4 
2 

4 
2 
4 

35.45 
35.45 
106.35 
106.35 

35.45 
70.9 

'32 

16 
160 

16 

16 
32 

31 
31 
124 
31 

200 
200 

HCI 

C102 
PCL 
POC13 

5fr 

H26 
CH4 
C2H2 
C2H4 
CH20 

C2H40, 

HgCl 
HgCl2 

Chlorine  dioxide    

Phosphorus  trichloride    .... 
Phosphorus  oxychloride  .... 
Phosphoric  anhydride      .... 
Phosphine          

Water  

Methane  ......... 

Ethylene  

Formaldehyde   ...          ... 

Mercuric  chloride 

oxychloride,  for  example,  the  quantity  was  106.35,  or  3  x  35.45. 
Hence  35.45  g.  can  be  taken  as  the  unit  quantity,  the  atomic  -weight 
of  the  element  chlorine.  In  terms  of  the  hypothesis,  this  is  the 
relative  weight  of  an  atom  of  chlorine,  as  compared  with  the  weight 


ATOMIC   WEIGHTS  199 

of  a  molecule  of  oxygen,  when  the  value  32  is  assigned  to  the 
latter.* 

In  the  preceding  table  a  few  sample  results  of  the  process  just  out- 
lined are  given.  The  first  column  contains  the  molar  weight,  i.e.  the 
weight  of  the  substance  which  occupies  the  G.M.V.  cube.  In  the 
other  columns  are  entered  the  weights  of  the  various  elements  which 
together  make  up  the  total  molar  weight.  To  simplify  the  numbers, 
the  value  1  is  used  for  hydrogen,  instead  of  1.008. 

To  contain  similar  data  for  all  the  volatile  compounds  of  every 
known  element,  a  huge  table,  of  which  this  might  be  a  small  corner, 
would  be  required.  With  such  a  table  at  hand  the  atomic  weight 
of  each  element  could  promptly  be  picked  out.  Thus,  in  the  carbon 
column  it  would  be  found  that  all  the  weights  of  carbon  were  either 
12  or  integral  multiples  of  12,  and  this  is  therefore  the  atomic  weight 
of  carbon.  Similarly  the  atomic  weight  of  oxygen  is  16,t  of  phos- 
phorus 31,  of  mercury  200  (see  Exercise  4). 

When  the  atomic  weights  have  finally  been  selected,  we  can  go 
through  the  table  and  change  all  the  numbers  into  multiples  of  the 
chosen  atomic  weights.  Thus,  for  70.9  we  write  2  x  35.45,  and  for 
106.35  we  write  3  X  35.45,  and  so  forth.  The  reader  can  prepare 
such  a  modification  of  the  table.  With  this  new  form  of  the  table 
before  us,  we  can,  finally,  replace  the  atomic  weights  by  the  symbols 
which  stand  for  them,  writing,  for  35.45,  Cl,  for  2  X  35.45,  Clj,  and 
so  forth.  The  results  of  doing  this  in  each  line,  i.e.  for  each  sub- 
stance, are  collected  at  the  ends  of  the  lines  in  the  last  column  of 
the  table.  The  reader  should  himself  repeat  the  substitutions  of  the 
symbols,  and  so  verify  the  formulae  given.  These  formulae,  since 
they  are  based  on  the  molecular  weights,  in  such  a  way  that  when 
the  numerical  values  are  substituted  for  the  symbols  the  total  restores 
to  us  the  molecular  weight,  are  called  molecular  formulae. 

It  will  now  be  seen  why  the  equivalents  (pp.  50, 51)  were  multiplied 
by  various  integers  in  making  the  chemical  units.  The  equivalent 
of  carbon  was  3.  That  is  to  say,  carbon  and  oxygen  combine  in  the 
ratio  3:8  (in  carbon  dioxide),  and  carbon  and  hydrogen  in  the 

*  It  should  be  noted  that  there  is  another  unit  quantity  of  chlorine,  namely  the 
molecular  weight,  or  weight  of  the  G.M.V.  of  the  substance.  This  is  the  unit 
quantity  of  free  chlorine.  But  we  are  dealing  now  with  compounds,  and  propor- 
tions in  combination,  so  that  free,  uncombined  chlorine,  and  other  elements  in 
free  condition  do  not  interest  us  at  present,  and  will  be  taken  up  later. 

t  The  difference  between  the  unit  quantity  of  oxygen  in  compounds  (namely, 
16)  and  the  unit  quantity  of  free  oxygen  (32)  will  be  discussed  presently. 


200  INORGANIC   CHEMISTRY 

ratio  3:1.008  (in  methane).  But  there  is  no  compound  of  carbon 
whose  molecular  weight  contains  less  than  12  parts  of  the  element. 
It  would  thus  lead  to  needless  complication  to  take  3  as  the  unit 
amount  of  carbon,  for  every  molecule  would  then  contain  four  units, 
or  some  multiple  of  four,  and  every  formula  C4  or  some  multiple  of 
C4.  We  choose  the  largest  units  of  combining  weight  that  we  can, 
in  order  that  the  coefficients  may  be  the  smallest  possible,  and  the 
resulting  formulae  the  simplest  possible.  Naturally  the  actual  ratios 
remain  the  same.  Thus,  for  carbon  dioxide  the  ratio  3  :  8  is 
replaced  by  12 : 32,  or  12 : 2  X  16,  or  C  :  20,  which  has  the  same 
value. 

Asa  definition,  the  atomic  weight  of  an  element  may  be  stated  to 
be:  The  smallest  of  the  weights  of  the  element  found  in  the  molecular 
weights  of  all  its  volatile  compounds,  so  far  as  these  have  been  exam- 
ined. Since  the  atomic  weight  is  always  a  multiple  of  the  equivalent 
weight  (by  unity,  or  some  other  integer),  it  might  also  be  denned  as : 
The  largest  integral  multiple  of  the  equivalent  which  can  be  contained 
in  the  molecular  weights  of  all  the  volatile  compounds  of  the  element. 
The  complete  list  of  accepted  atomic  weights  is  printed  on  the  inside 
of  the  cover  at  the  back  of  this  book. 

Advantages  of  Atomic  Weights  over  Equivalents.  —  Since  the 
method  of  determining  atomic  weights  depends  on  rather  complex 
reasoning,  and  involves  much  experimental  work,  the  question  may  be 
asked  whether,  when  found,  they  are  worth  all  the  trouble.  It  is 
manifest  that  equivalents  are  much  simpler  in  nature,  and  much  more 
easily  ascertained  than  atomic  weights.  It  will  be  expected,  there- 
fore, that  we  shall  be  able  to  show  that  the  units  Na  =  23,  Cu  =  63.6, 
Al  =  27.1,  C  =  12,  etc.,  give  a  better  view  of  the  relations  of  the  ele- 
ments than  do  the  equivalents  23,  31.8,  9.03,  and  3,  respectively. 
Now,  the  atomic  weights  are  as  good  as  equivalent  weights,  for  the 
purpose  of  acting  as  units,  in  terms  of  which  to  express  combining 
proportions,  and  possess,  besides,  several  (at  least  five)  important 
properties  or  uses  which  equivalent  weights  entirely  lack. 

Of  these  valuable  properties,  the  first  two  have  been  mentioned 
already : 

1.  Being  very  often  themselves  larger  numbers  than  the  equiva- 
lents, atomic  weights  are  often  multiplied  by  smaller  coefficients  (see 
above).     This  simplifies  our  equations. 

2.  The  atomic  weight  of  an  element  can  have  but  one  value,  and 


ATOMIC    WEIGHTS 


201 


is  definitely  determinable.  Most  equivalents  have  more  than  one 
value,  because  an  element,  when  it  gives  several  series  of  compounds 
(p»  105),  will  have  as  many  different  equivalents. 

3.  The  atomic  weight  of  an  element  has  a  valence  (p.  101),  while 
equivalents  are  equi-valent.     While   valence  is  a  helpful  conception 
in  all  branches  of  chemistry,  organic  chemistry  is  especially  indebted 
to  the  conception  of  the  quadrivalence  of  carbon  for  much  of  its  devel- 
opment and  most  of  its  organization.     The  full  illustration  of  this 
point  is  beyond  the  limits  of  the  present  book. 

4.  The  periodic  system  (q.v,),  the  basis  of  a  plan  for  classifying 
the  properties  of  all  chemical  substances,  is  founded  upon  the  atomic 
weights. 

5.  Dulong  and  Petit's  law  is  based  upon  atomic  weights.     This 
law  furnishes  also  an  alternative  means  of  determining  atomic  weights 
that  has   frequently  rendered    valuable  service,  and  on  this  account 
forms  the  subject  of  the  next  section. 

Dulong  and  Petit  >s  Law,  an  Alternative  Means  of  Deter- 
mining  Atomic  Weights. — It  was  first  pointed  out  (1818)  by  Dulong 
and  Petit,  of  the  Ecole  Polytechnique  in  Paris,  that  when  the  atomic 
weights  of  the  elements  were  multiplied  by  the  specific  heats  of  the 
simple  substances  in  the  solid  condition,  the  products  •were  approxi- 
mately the  same  in  all  cases.  In  other  words,  the  specific  heats  are 
inversely  proportional  to  the  magnitudes  of  the  atomic  weights.  The 
table,  in  which  round  numbers  have  been  used  for  the  atomic  weights, 
shows  that  the  product  lies  usually  between  6  and  7,  averaging 
about  60  4 : 


Element. 

Atomic 
Wt. 

Sp.  Ht. 

Pro- 
duct. 

Element. 

Atomic 
Wt. 

Sp.  Ht. 

Pro- 
duct. 

Lithium      .     .     . 

7 

.94 

6.6 

Iron      .... 

56 

.112 

6.3 

Sodium  .     .     . 

23 

.29 

6.7 

Zinc     .... 

65.4 

.093 

6.1 

Magnesium 

24.4 

.245 

6.0 

Bromine  (Solid) 

80 

.084 

6.7 

Silicon   .... 

28.4 

.16 

4.5 

Gold     .... 

197 

.032 

6.3 

Phosphorus 

81 

.19 

5.9 

Mercury 

200 

.0335 

6.7 

(Yellow) 

(Solid) 

Calcium      .     .     . 

40 

,170 

6.8 

Uranium  .     .     . 

238.5 

.0276 

6.6 

Another  way  of  expressing  this  law  will  give  it  greater  chemical 
significance.     The  specific  heats  are  the  amounts  of  heat  required  to 


202  INORGANIC   CHEMISTRY 

raise  equal  weights  of  the  various  elements  through  one  degree.  Now 
these  equal  weights  contain  fewer  chemical  units  in  proportion  as  the 
chemical  unit  weight  is  greater.  Hence  this  law  may  be  put  in  the 
form:  Equal  amounts  of  heat  will  raise  atomic  weights  of  all  elements 
through  equal  intervals  of  temperature. 

This  being  true,  the  equivalents,  if  used  instead  of  the  atomic 
weights,  must  give  widely  varying  products.  The  quantities  of  heat 
required  to  raise  equivalent  weights  through  one  degree  are  either 
equal  to,  or  are  fractions  of,  those  required  for  the  atomic  weights, 
according  to  the  valence  of  the  element.  Hence  the  law  applies 
only  to  atomic  weights,  and  not  to  equivalents. 

It  will  be  seen  at  once  that  although  the  law  of  Dulong  and  Petit 
is  purely  empirical,  it  may  nevertheless  be  used  for  fixing  the  atomic 
weight  of  an  element  of  which  no  volatile  compounds  are  known.  We 
can  always  measure  the  equivalent  with  considerable  exactness,  and, 
when  this  has  been  multiplied  by  the  specific  heat  of  the  free  sub- 
stance, we  can  see  at  a  glance  what  integral  factor  will  raise  the 
product  to  the  neighborhood  of  6.4.  For  example,  analysis  shows  us 
that  in  calcium  chloride  the  proportion  of  chlorine  to  calcium,  using 
the  known  atomic  weight  of  chlorine  as  one  term  of  the  proportion, 
is  35.5  :  20.  If  calcium  is  univalent,  20  is  its  atomic  weight.  If  it 
is  bivalent,  two  units  of  chlorine  are  combined  with  40  parts  of 
calcium,  and  40  is  its  atomic  weight.  If  it  is  trivalent,  three  units  of 
chlorine  are  united  with  60  parts  of  calcium,  etc.  All  we  learn  in 
reference  to  the  atomic  weight  of  calcium  from  this  analysis  is  that 
its  value  is  20  or  some  integral  multiple  of  20.  Nor  can  we  fix  the 
upper  limit,  for  we  are  unable  to  obtain  the  weight  of  a  known 
volume  of  calcium  chloride  vapor  and  so  determine  the  molecular 
weight.  But  the  specific  heat  of  solid  calcium  being  0.170,  we  mul- 
tiply this  number  by  20,  and  get  the  product  3.4.  This  is  only  half 
large  enough,  so  we  assume  that  40  is  a  more  probable  value  for  the 
atomic  weight  of  calcium.  The  product  is  then  6.8,  which  agrees 
fairly  well  with  the  average  for  other  elements.  We  decide,  there- 
fore, that  the  symbol  Ca  shall  represent  forty  parts  by  weight.  The 
formula  of  calcium  chloride  is  therefore  CaCl2,  and  calcium  is  bivalent. 

It  will  be  seen  that  this  does  not  supply  us  with  a  method  of  ascertaining 
chemical  unit  weights  independently  of  any  chemical  experiment.  We  cannot 
measure  the  specific  heat  and  use  the  quotient  from  division  of  this  number  into 
6.4,  for  we  do  not  know  in  advance  that  the  product  for  the  element  will  have 
exactly  this  value.  It  may  be  below  6,  or  it  may  be  as  high  as  7.  In  the  case  of 


MOLECULAR   EQUATIONS  203 

calcium,  for  example,  6.4  -f-  0.17  =  37.65.  Now  37.65  is  5  per  cent  below  the  real 
value  of  the  chemical  unit,  and  even  the  roughest  measurement  of  a  chemical 
combining  weight  need  never  be  more  than  1  per  cent  in  error.  Hence  the  atomic 
weight  must  be  founded  upon  the  determination  of  the  equivalent,  which  can 
be  measured  with  accuracy.  The  rule  discussed  in  this  section  can  be  used  only 
to  ascertain  what  multiple  of  the  equivalent  shall  be  accepted  as  the  atomic  weight 
after  the  equivalent  itself  has  been  measured  with  care.  In  other  words,  this  is  a 
method  of  adjusting  the  result  of  chemical  experimentation,  and  cannot  supersede 
it  altogether. 

The  existence  of  the  law  of  Dulong  and  Petit  and  the  periodic  law,  together 
with  the  services  of  structural  formulae  to  organic  chemistry,  all  demonstrate  that 
atomic  weights  are  of  vastly  greater  significance  in  the  science  than  are  equivalent 
weights.  And  there  are  other  immense  ranges  of  facts,  aside  from  those  covered 
by  these  conceptions,  which  are  all  dependent  upon  the  atomic  weights.  That 
almost  the  whole  systematization  that  has  been  secured  in  chemistry  should  thus 
center  in  this  one  point,  furnishes  the  strongest  circumstantial  evidence  that 
Avogadro's  hypothesis  represents  very  closely  some  fundamental  property  of  all 
gases.  This  independent  inductive  evidence  in  favor  of  Avogadro's  principle  is 
especially  worth  noting  because  the  deduction  of  the  principle  from  the  data  of  the 
kinetic-molecular  hypothesis  is  not  absolutely  rigid.  It  involves  certain  assump- 
tions which,  while  they  are  plausible  enough,  are  still  assumptions. 

MOLECULAR  FORMULA. 

Molecular  Formulce  of  Compounds.  —  If  the  molar  formulae 
in  the  table  (p.  198)  be  examined  it  will  be  observed  that  several  are 
not  in  their  simplest  terms.  Thus,  the  formula  of  acetylene  is  C2H2. 
The  formula  CH  would  represent  the  composition  of  the  substance 
equally  well,  for  12  :  1  is  the  same  as  24  : 2.  But  the  formula  CH 
gives  a  total  of  only  13,  while  C2H2  shows  the  total  weight  of  the 
molecule  to  be  26  and  records  for  us  therefore  the  weight  of  the  G.M.  V., 
as  well  as  the  composition  of  the  substance.  And  we  shall  find  this 
additional  property,  peculiar  to  the  molecular  formula,  to  be  a  feature 
of  the  greatest  practical  value.  Some  of  the  practical  uses  of  this 
improvement  in  our  formula  will  be  illustrated  in  this  chapter,  and 
there  is  an  example  of  one  of  them  in  the  table  itself.  Thus,  the 
molecular  formula  of  acetic  acid  is  C2H402,  and  not  the  simpler, 
identical  proportion  CH20.  The  latter  is  the  molecular  formula  of 
a  totally  different  substance,  formaldehyde,  now  much  used  as  a  dis- 
infectant. The  vapor  of  this  substance  has  only  half  the  density  of 
acetic  acid  vapor,  and  this  fact,  recorded  in  the  formula,  helps  to  re- 
mind us  that  the  substances  are  different.  Still  another  substance  of 
the  same  composition  is  grape  sugar  (dextrose),  C6H1006  (see  Exercise 
12).  In  addition  to  this  and  other  practical  advantages,  molecular 


204 


'INORGANIC   CHEMISTRY 


formulae  satisfy  also  the  claim  of  logical  consistence.  If  the  sym- 
bols represent  the  atomic  weights,  the  formulae  should  be  constructed 
so  as  to  represent  the  molecular  weights. 

Molecular  formulae  like  C2H2  and  C2H402  are  easily  interpreted 
in  terms  of  the  atomic  hypothesis.  C  represents  one  atom  of  carbon 
and  H,  one  atom  of  hydrogen.  But  there  is  no  reason  why  a  mole- 
cule of  acetylene  should  not  contain  two  atoms  of  each  kind.  Simi- 
larly, the  molecule  of  formaldehyde  contains  four  atoms  (CH20),  and 
one  of  acetic  acid  eight  atoms  (C2H4O2),  and  one  of  dextrose  twenty- 
four  atoms  (C6H1206),  although  the  relative  numbers  of  each  kind  are 
the  same.  Indeed,  this  hypothesis  helps  to  clear  the  matter  up,  for 
chemists  go  so  far  as  to  account  for  the  chemical  behavior  of  the 
substances  by  an  imagined  geometrical  arrangement  of  the  atoms  in 
their  molecules,  and  these  three  kinds  of  molecules  are  supposed  to 
differ  in  structure  as  well  as  in  the  number  of  atoms  they  contain. 

The  Molecular  Weights  and  Formulce  of  Elementary  Sub- 
stances.—  The  following  table  gives  the  densities  of  some  elementary 
substances,  including  those  of  which  the  substances  last  discussed  are 
compounds.  The  first  column  shows  the  atomic  weight,  which  in  each 
case  is  the  minimum  weight  of  the  element  found  in  a  G.M.V.  of 


Atomic 

Sym- 

Density, 

Density  Fac- 

Formula 

* 

Weight. 

bol. 

O  =  32. 

to  rized. 

Element. 

Oxygen   

16.00 

O 

32.00 

2X16.00 

°2 

Hydrogen  .... 

1  008 

H 

2  016 

2x1  008 

H, 

Chlorine     

35.46 

Cl 

70.90 

2x35.45 

ct 

Phosphorus    

31  0 

P 

124  0 

4x31  0 

P/ 

Mercury 

200  0 

Hff 

200  0 

1x200  0 

Hg 

Ozone      

16.00 

0 

48  00 

3x16.00 

03 

Cadmium    

112.4 

Cd 

112.4 

1x112.4 

Cd 

Potassium  

39.15 

K 

39  15 

1x39.15 

K 

Sodium   

23.05 

Na 

23.05 

1x23.05 

Na 

Zinc.               .        ... 

65  4 

Zn 

65  4 

1x65.4 

Zn 

any  compound.  For  example,  16  g.  of  oxygen  and  35.45  g.  of  chlor- 
ine are  the  weights  in  the  amounts  of  water  vapor  and  hydrogen 
chloride,  respectively,  which  fill  the  cube  (22.4  liters).  The  symbol, 
in  the  next  column,  stands  for  this  quantity  and  occurs  in  many 
formulae,  such  as  HO  and  HC1.  It  represents  the  combining  unit  or 


MOLECULAR  EQUATIONS  205 

atom.  In  the  third  column  is  given  the  density  of  the  free,  elemen- 
tary substance.  This  number  of  grams  of  the  simple  substance  fills 
the  Gr.M.V.  and  this  number  is  the  molecular  weight.  It  shows  the 
weight  of  the  molecule  relative  to  the  weights  of  the  other  molecules  in 
the  same  column,  and  to  the  weights  of  the  atoms  in  the  first  column. 
In  the  last  two  columns  are  given  the  densities  resolved  into  multiples 
of  the  atomic  weights  and  the  corresponding  formulae. 

The  reader  cannot  fail  to  note  a  striking  peculiarity.  In  the  case 
of  chlorine  the  molecular  weight  is  70.9,  while  the  atomic  weight  is 
35.45.  With  hydrogen  and  oxygen,  also,  the  molecular  weight  con- 
tains two  atomic  weights.  Yet  this  is  not  a  general  rule,  for  with 
mercury  and  several  other  elements  the  molecular  and  atomic  weights 
are  alike,  while  with  phosphorus  the  molecular  is  four  times  the 
atomic  weight.  Evidently  there  is  no  rule,  and  each  element  has  to  be 
subjected  to  separate  experimental  study.  The  result  is  that  for  free, 
elementary  chlorine  we  use  the  molecular  formula  C19,  for  free  hydro- 
gen H2,  for  elementary,  uncombined  oxygen  the  formula  02.  For  a 
substance  like  phosphorus,  which  is  not  a  gas  and  is  not  often  measured 
as  a  vapor,  the  formula  P  is  commonly,  employed  by  chemists,  to  avoid 
the  larger  coefficients  which  P4  introduces  into  equations,  although 
theoretically  the  latter  formula  would  be  the  strictly  correct  one. 

The  case  of  oxygen  demonstrates  clearly  the  necessity  of  using 
molecular  formulas,  even  for  simple  substances.  The  table  shows  two 
substances  containing  nothing  but  oxygen.  Ozone  (q.v.)  has  a  molec- 
ular weight  48,  being  a  gas  exactly  one-half  heavier  than  ordinary 
oxygen.  Its  formula,  therefore,  is  03,  while  that  of  oxygen  is  O2. 
Oxygen  and  ozone  are  entirely  different  chemical  individuals.  The 
latter  has,  for  example,  a  strong  odor  and  is  much  more  active.  Thus 
polished  silver  remains  bright  indefinitely  in  pure  oxygen,  but  oxidizes 
quickly  when  placed  in  ozone. 

To  avoid  a  common  error,  the  reader  should  note  that  to  learn  the 
atomic  weight  of  an  element,  we  do  not  measure  the  molecular  weight 
of  the  simple  substance.  The  molecular  weight  of  the  elementary 
substance  may  be  a  multiple  of  the  atomic  weight,  and  we  find  out 
whether  it  is  such  a  multiple  only  after  the  atomic  weight  has  been 
determined.  The  atomic  weight  is  the  unit  weight  used  in  compounds, 
and  can  be  ascertained  only  by  a  study  of  compounds.  The  molecular 
weight  of  the  free  element  gives  us  only  a  value  which  we  know  must 
be  a  multiple  of  the  atomic  weight,  by  1  or  some  other  integer.  Mol. 
Wt.  =  At.  Wt.  x  x,  where  x  is  1  or  some  other  integer. 


206 


INORGANIC   CHEMISTRY 


Further  Discussion  of  the  Molecular  Formulce  of  Elemen- 
tary Substances.  —  Some  further  explanation  may  be  required,  to 
the  end  that  the  reader  may  be  reconciled  to  accepting  the  formulae 
CLj,  O2,  and  so  forth.  In  the  first  place,  he  should  note  how  these 
formulae  arose.  If  we  accept  Avogadro's  hypothesis,  and  the  inference 
from  it  to  the  effect  that  the  densities  of  gases  are  in  the  same  ratio 
as  the  weights  of  their  individual  molecules,  then  we  cannot  escape 
the  conclusion  to  which  measuring  the  relative  densities  of  free 
chlorine  and  hydrogen  chloride,  for  example,  leads.  The  ratio  of 
their  densities  is  70.9 :  36.45.  That  is  to  say,  the  relative  weights 
of  a  molecule  of  chlorine  and  a  molecule  of  hydrogen  chloride  stand 
in  this  ratio.  The  molecule  of  chlorine  is  nearly  twice  as  heavy  as 
the  molecule  of  the  compound,  and  there  cannot  therefore  be  a  whole 
molecule  of  chlorine  in  a  molecule  of  hydrogen  chloride.  In  fact,  we 
perceive  at  once  that  the  molecule  of  hydrogen  chloride  must  contain 
only  half  a  molecule  of  chlorine  (35.45),  together  with  half  a  molecule 
of  hydrogen  (1).  In  other  words,  if  the  molecule  of  free  chlorine 
were  to  be  taken  as  the  atom  of  the  element,  then  the  molecule  of 
hydrogen  chloride  would  contain  only  half  an  atom  of  chlorine, 
which  would  be  contrary  to  our  decision  to  take  as  atoms  quantities 
which  are  not  divided.  So  we  choose  the  other  horn  of  the  dilemma, 
and  say  that  the  specimen  of  chlorine  in  the  molecule  of  hydrogen 
chloride  is  a  whole  atom  and  that  therefore  the  amount  of  chlorine 
in  the  molecule  of  free  chlorine  is  two  atoms,  and  its  formula  Cl,. 
Similarly,  the  weight  of  hydrogen  in  the  molecule  of  hydrogen  chlo- 
ride is  1.008,  while  that  of  the  molecule  of  hydrogen  is  2.016,  so  that 
there  are  two  atoms  in  the  molecule  of  free  hydrogen  and  its  formula 
is  Hj  Reasoning  in  like  manner  from  the  molecular  weights  of 
oxygen  (32)  and  water  (18)  we  reach  the  conclusion  that  the  molecule 
of  oxygen  is  diatomic  (02). 

Still  another  way  of  looking  at  the  same  facts  may  shed  light  on 
the  matter.  When  hydrogen  and  chlorine  combine,  one  volume  of 
each  of  these  gases  gives  two  volumes  of  hydrogen  chloride  (p.  122). 
Let  us  imagine  the  experiment  to  be  made  with  minute  volumes 
holding  one  hundred  molecules  each  : 

HYDROGEN  CHLORIDE  HYDROGEN        CHLORINE 

came  from 

The  200  molecules  of  hydrogen  chloride  must  contain  at  least  200 
fragments  of  chlorine,  since  there  is  a  sample  in  each  molecule.     Now 


100 

100 

APPLICATIONS   OF   MOLECULAR   EQUATIONS  207 

the  200  fragments  of  chlorine  came  from  a  volume  containing  only 
100  molecules  of  chlorine.  Each  of  these  must  therefore  have  been 
split  in  the  chemical  action.  Hence  the  molecules  of  free  chlorine 
contain  at  least  two  atoms.  Parallel  reasoning  leads  to  the  conclusion 
that  the  molecules  of  free  hydrogen  are  likewise  diatomic.  If  we 
consider  the  molecular  formula  of  a  substance  as  representing  one 
molecule  (see  below),  the  equation  for  this  action  is : 

H2  +  C12->2HC1. 

There  are  two  molecules  on  each  side  of  the  equation,  and  this  corre- 
sponds with  the  fact  that  there  is  no  change  in  the  total  volume. 
Again,  we  find  that  one  volume  of  oxygen  furnishes  enough  of  the 
element  for  two  volumes  of  water  vapor  (p.  125).  We  infer  therefore 
that  each  molecule  of  .oxygen  is  divided  into  two  parts  in  the  action. 
And  in  like  manner,  when  we  find  that  one  volume  of  phosphorus 
vapor,  in  combination  with  six  volumes  of  chlorine,  gives  four 
volumes  of  phosphorus  trichloride  vapor,  we  infer  that  every  molecule 
of  phosphorus  furnished  enough  of  the  element  for  four  molecules  of 
phosphorus  trichloride,  and  contained  therefore  four  atoms  (see  Exer- 
cise 7). 

P4  +  6CL,  -4  4PC13. 

The  simple  fact  that  hydrogen  and  oxygen,  when  mixed,  do  not 
combine  (p.  109)  may  assist  in  reconciling  us  to  the  diatomic  nature 
of  their  molecules.  Some  part  of  the  mixture  has  to  be  heated 
strongly  to  start  the  interaction.  Now  the  molecular  formulae,  H0 
and  02,  suggest  that  each  gas  is  really  in  combination  already  (with 
itself),  and  they  therefore  explain  to  some  extent  the  indifference  of 
the  gases  towards  one  another.  If  the  molecules  were  free  atoms,  they 
could  not  encounter  one  another  continually  as  they  move  about,  and 
yet  escape  combination  as  we  observe  that  they  do.  We  may  imagine 
that  the  primary  effect  of  heating  is  to  decompose  some  of  the  mole- 
cules, and  liberate  hydrogen  and  oxygen  in  the  atomic  condition,  and 
that  the  combination  of  these  atoms  starts  the  explosion  of  the  whole 
mass.  Of  course  this  explanation  is  based  upon  our  hypothesis  and,  as 
such,  is  of  the  same  imaginary  description  as  everything  connected 
with  that  hypothesis.  It  cannot  be  verified  by  experiment. 

APPLICATIONS. 

Applications:  Interactions  Between  Gases.  —  According  to 
Avogadro's  hypothesis,  if  we  filled  a  succession  of  vessels  of  equal 


208  INORGANIC   CHEMISTRY 

dimensions  with  different  gases,  and  could  arrest  the  motion  of  the 
particles  and  observe  their  disposition,  we  should  find  that  the  average 
distance  from  particle  to  particle  would  be  the  same  in  all  cases.  This 
would  be  true  whether  our  vessels  were  filled  with  single  gases,  with 
homogeneous  mixtures,  or  with  gases  in  layers.  Such  being  the  case, 
if  any  chemical  change  is  brought  about  in  the  mass  which  results  in  a 
multiplication  of  the  molecules,  it  is  evident  that  the  volume  will  have 
to  increase  in  order  that  the  spacing  may  remain  the  same  as  before. 
If  any  chemical  action  results  in  a  diminution  of  the  number  of  mole- 
cules, then  a  shrinkage  must  take  place  in  order  that  the  spacing  may 
be  preserved  as  before.  Thus,  in  a  mixture  of  hydrogen  and  chlorine, 
according  to  our  hypothesis,  when  interaction  to  produce  hydrogen 
chloride  occurs,  neighboring  molecules  of  hydrogen  and  chlorine  simply 
exchange  units,  so  that  HH  +  C1C1  becomes  HC1  +  C1H.  There 
being  no  alteration  in  the  number  of  particles,  no  change  in  volume 
occurs.  In  the  case  of  water,  on  the  other  hand, 

HH  +  00  +  HH     becomes     HOH  +  HOH. 

Since  the  oxygen  molecules,  which  form  a  third  of  the  whole, 
disappear  into  the  molecules  of  hydrogen,  the  tendency  to  pre- 
serve spacing  results  in  a  diminution  of  the  volume  by  one-third 
(p.  125). 

This  method  of  looking  upon  chemical  interactions  between  gases 
gives  us  the  nearest  sight  which  we  can  have  of  the  behavior  of  the 
molecules  themselves.  We  cannot  perceive  the  individual  molecules, 
but,  in  consequence  of  the  spatial  arrangement  which  we  suppose  them 
to  observe,  the  change  in  the  whole  volume  of  a  large  aggregate  of 
molecules  enables  us  to  draw  conclusions  at  once  in  regard  to  the  be- 
havior of  the  single  molecules  in  detail. 

Applications :  Molecular  Equations.  —  To  utilize  the  foregoing 
considerations,  chemists  always  employ  in  their  equations  the  molec- 
ular formulae  for  the  gases  and  easily  vaporized  substances  concerned. 
Thus  far,  we  have  used  the  equation : 

H   +   Cl  -»  HC1, 
WEIGHTS:  1.008       36.45      36.458 

and  the  information  it  contained  was  exhausted  when  we  had  placed 
below  the  symbols  the  weights  for  which  they  stood.  But  the  molec- 


APPLICATIONS   OF   MOLECULAR   EQUATIONS  209 

ular  equation  is  much  more  instructive.     The  following  shows  the  in- 
terpretations to  which  the  molecular  equation  is  subject: 

H2      +  C12    -4     2HC1. 

WEIGHTS  :                    2.016  g.  70.90  g.       2  x  36.458  (=  72.916)  g. 

VOLUMES:                   22.41.  22.41.          2x22.41. 

MOLECULES  :                     1  1                       2 

The  weights,  although  doubled,  show  the  same  proportions,  so  that 
questions  of  weight  are  answered  as  easily  as  before.  These  weights, 
however,  being  molecular  weights,  or  multiples  thereof,  can  be  trans- 
lated at  once  into  volumes,  and  questions  about  volumes  can  also  be 
answered.  Finally  the  relative  numbers  of  each  kind  of  molecules 
can  be  read  from  this  equation,  for  the  coefficients  in  front  of  the 
formulae  represent  these  numbers.  Where  no  coefficient  is  written,  1  is 
to  be  understood.  The  application  of  these  properties  of  molecular 
equations  is  illustrated  below.  Before  applying  these  equations,  how- 
ever, we  must  first  learn  how  to  make  them. 

To  make  a  molecular  equation,  we  first  make  an  equation  according 
to  the  rules  already  explained  (p.  57) .  An  equation  like  that  given  for 
the  interaction  of  oxygen  with  hydrogen  chloride  (Deacon's  process, 
p.  170) :  2HC1  +  0  -»  H20  +  2C1,  is  the  result.  Then  we  adjust 
the  equation  so  that  molecular  formulae  are  used  throughout.  2C1 
becomes  at  once  C12.  The  oxygen,  however,  must  also  appear  as  02, 
or  a  multiple  of  this,  in  such  equations.  Hence  the  whole  equation 
must  be  multiplied  by  2 : 

4HC1  +  02  ->  2H20  +  2C12. 

Again,  the  equation  for  the  preparation  of  chlorine  from  potassium 
permanganate  and  hydrochloric  acid  (p.  173)  becomes : 

2KMn04  +  16HC1  ->  8H2O  +  2KC1  +  2MnCl2  +  5C12. 

Every  equation  containing  an  odd  number  of  atoms  of  a  substance 
whose  molecules  are  diatomic  must  be  multiplied  by  2.  Again,  mer- 
curic oxide  decomposes  to  give  mercury  vapor  and  oxygen  (p.  59),  and 
the  molecules  of  mercury  are  monatomic  and  those  of  oxygen  diatomic, 
so  we  write  : 

2HgO  -+  2Hg  +  02. 

Finally,  the  formulae  of  substances  which  are  solid  or  liquid,  and  can- 
not be  easily  vaporized,  are  written  in  the  simplest  terms.  Thus,  since 
substances  like  the  copper  in  the  following  equation  are  involatile,  the 


210  INORGANIC   CHEMISTRY 

molecular  weights  of  such  substances  are  unknown,  and  their  molec- 
ular formulae  likewise :  2Cu  +  02  — *  2CuO.  Furthermore,  in  the 
case  of  substances  which  can  be  volatilized,  although  the  molecular 
weights  and  molecular  formulae  may  therefore  be  known,  we  do  not 
usually  employ  the  molecular  formulae  if  the  substance  is  not  used  in 
the  form  of  vapor  in  the  laboratory.  Thus,  the  molecular  formula  of 
phosphoric  anhydride  is  P4O10  (p.  198).  But  we  generally  make,  and 
use,  only  the  solid  form,  and  not  the  vapor,  in  actual  work.  Hence 
the  action  with  water  is  usually  written  as  we  have  given  it  (p.  71), 
rather  than  in  the  form  :  P4O10  +  6H2O  -»  4H3P04. 

The  applications  of  the  properties  of  molecular  equations  (which 
will  be  used  exclusively  hereafter)  may  now  be  illustrated  in  detail. 

Applications  :  To  Arithmetical  Problems.  —  1.  When  a  prob- 
lem in  regard  to  weights  of  material  used  or  produced  in  a  given  action 
is  to  be  solved,  the  molecular  equation  is  to  be  written  and  the  weights 
inserted  beneath  the  formulae.  The  mode  of  calculation  has  been 
described  already  (p.  59). 

2.  When  a  problem  involving  weights  and  volumes  is  to  be 
solved,  the  molecular  equation  is  to  be  written,  and  both  the  weights 
and  volumes  are  to  be  inserted.  Note,  however,  that  only  the  volumes 
of  the  substances  in  the  gaseous  condition  are  considered. 

For  example,  what  volume  of  oxygen  is  obtained  from  60  g.  of 
potassium  chlorate?  The  molecular  equation,  made  as  shown  above, 
(p.  209),  together  with  the  full  interpretation,  are  as  follows  : 

2KC103          ->  2KC1      +      302. 

J2  (39.15  +  35.45  +48)        2(39.15  +  35.45)  3_X_32 

WEIGHTS:    j  245.2  g.  149.2  g.  ^96gT/ 

VOLUMES  ;  3  X  22.4 1. 

(Observe  that  no  volumes  are  given  under  the  chlorate  and  chloride 
of  potassium.  This  is  because  their  volumes  in  the  gaseous  condition 
can  be  of  no  practical  use,  since  they  are  solids  which  are  melted,  but 
not  vaporized  during  this,  or  any  action  in  which  we  employ  them.) 
Now,  as  to  the  problem  in  hand,  it  is  concerned  with  a  weight  of  potas- 
sium chlorate  and  a  volume  of  oxygen.  Reading  from  the  equation, 
our  information  on  these  points  is  that  245.2  g.  of  potassium  chlorate 
give  67.2  liters  (observe  that  the  coefficients  are  used,  as  well  as  the 
molecular  weights,  in  these  numbers)  of  oxygen  at  0°  and  760  mm., 
and  the  question  is,  What  volume  will  60  g.  give?  By  proportion, 


APPLICATIONS   OF   MOLECULAR    EQUATIONS  211 

245.2  g. :  67.21:  :  60  g.  :  x  1.,  where  x  =  16.4  liters.  If  a  different 
temperature  and  pressure  had  been  specified,  either  the  volume  in  the 
equation,  or  the  answer,  would  have  had  to  be  converted  by  rule  to 
the  given  conditions.  It  saves  time  not  to  write  out,  as  above,  the 
whole  interpretation,  but  only  the  parts  required.  For  example,  if  the 
question  is :  What  volume  of  chlorine  is  needed  to  give  25  g.  of  alu- 
minium chloride,  we  may,  if  we  choose,  omit  -all  the  data  excepting  the 
volume  of  the  chlorine  and  the  weight  of  the  aluminium  chloride,  thus  : 

2A1  +     3C1,     — »     2A1C13 
3x22.41.     2  x  133.45  g. 

The  volume  of  chlorine  required  is  3x22.4x25  -T- (2x133.45)  liters. 
These  illustrations  show  the  method  of  calculating  actual  volumes  (see 
Exercises  8,  9). 

3.  If  the  question  concerns  relative  volumes  only,  then  it  is 
simplest  to  use  the  interpretation  of  the  equation  in  terms  of  mole- 
cules. For  example,  What  relative  volumes  of  hydrogen  chloride  and 
oxygen  are  required  in  Deacon's  process  ?  The  molecular  equation  is 
(p.  209)  : 

4HC1  +  02  ->  2H20  +  2C12. 
MOLECULES :  4122 

Since  equal  numbers  of  molecules  of  gases  occupy  equal  volumes,  the 
proportion  4  molecules  of  hydrogen  chloride  to  1  molecule  of  oxygen 
shows  the  ratio  to  be  4 :  1  by  volume.  Similarly,  every  4  molecules 
of  hydrogen  chloride  give  2  molecules  of  chlorine,  so  that  the  ratio  of 
these  substances  by  volume  is  4 :  2,  or  2 :  1. 

In  regard  to  the  water,  since  that  is  not  a  gas  at  common  tempera- 
tures, the  question,  if  asked,  must  be  more  specific :  What  are  the 
relative  volumes  of  steam  and  chlorine  in  the  product,  as  commonly 
delivered  by  the  action  at  400°  ?  It  is  2  :  2,  or  1  :  1.  What  are  the 
relative  volumes  of  water  and  chlorine,  after  the  products  have  cooled 
to  room  temperature  ?  The  water  is  no  longer  a  gas,  so  that  it  occu- 
pies, relatively,  almost  no  volume.* 

What  is  the  total  volume-change  in  the  foregoing  action  above 

*  Of  course  if  an  exact  answer  must  be  given,  it  can  be  given.  But  for  this 
we  require  the  weight  and  specific  gravity  of  the  product.  Thus,  2H2O  represents 
2  x  18  g.  of  water.  The  sp.  gr.  of  water  is  1.  Therefore  the  volume  of  water 
formed  is  36  c.c.  The  volume  of  2C12  is  2  x  22.4,  or  44.8  liters  at  0°.  The  ratio 
of  water  to  chlorine  by  volume  at  0°  is  therefore  36  :  44,800.  But,  as  a  rule,  we 
simply  give  the  volumes  of  solids  and  liquids  as  zero,  compared  with  those  of  the 
gases  concerned  in  the  same  action. 


212  INORGANIC   CHEMISTRY 

100°?  It  is  a  change  from  5  molecules  to  4.  The  volume  changes  in 
the  same  ratio.  But  at  0°  the  volume-change  is  from  5  volumes  to  2, 
for  the  water  does  not  appreciably  add  to  the  volume  of  the  products 
(see  Exercises  10,  12). 

4.  When  we  know  the  molecular  formulae  of  the  single  substances 
concerned  in  an  action,  the  equation  can  be  made,  and  the  relative 
volumes   determined,  -without    actual    measurement.      For    example: 
What  volume-change   will   be   observed   when  a  mixture   of  carbon 
monoxide  and  oxygen  has  exploded,  and  the  temperature  has  once 
more  reached  that  of  the  room  ?     The  molecular  formulas  are  CO,  O9, 
and  C02.     The  equation  representing  the  weights  is  CO  -f  O  — >  C02. 
The  molecule  of  oxygen,  however,  being  O2,  we  cannot  employ  less 
than  this  quantity  in  a  molecular   equation,  so  that  the   equation 
becomes : 

2CO  +  02  ->  2C02. 

Three  molecules,  therefore,  give  two,  throughout  the  whole  mass,  and 
therefore  three  volumes  will  become  two,  if  the  pressure  and  tempera- 
ture are  the  same  at  the  beginning  and  end  of  the  action. 

If  we  remember  that  all  volatile  compounds  of  carbon  and  hydro- 
gen burn  to  form  water  and  carbon  dioxide,  the  molecular  equation  for 
any  such  combustion  may  easily  be  made,  and  the  volumes  of  all  the 
materials  ascertained.  When  water  is  a  product,  only  its  volume  as 
steam  is  given  by  the  equation  (see  Exercises  11,  12). 

5.  Knowing  by  heart  the  molecular  formulae  of  gaseous  substances, 
as  we  must  know  them  for  many  purposes,  it  is  unnecessary  to  burden 
our  minds  with  other  data  in  regard  to  the  relative  weights  of  gases. 
Is  hydrogen  chloride  (HC1)  heavier  or  lighter  than  carbon  dioxide 
(C02)  ?      These   formulas   represent   the   weights   of   equal  volumes 
(22.4  1.),  namely  36.45  g.  and  44  g.,  respectively.     Hence  the  former 
gas  is  a  little  lighter.     Remembering  that  the  G.M.V.  of  air  weighs 
28.955  g.,  we  can  compare  the  weight  of  any  gas  with  that  of  air  in 
the  same  way. 

What  are  the  relative  weights  of  acetylene  (C2H2,  p.  198)  and 
sulphur  dioxide  (S02)  as  compared  with  air  ?  The  G.M.V.  cube  holds 
formula-weights  of  the  first  two,  namely  26  g.  and  64  g.,  and  28.955  g. 
of  air.  Hence  acetylene  is  a  little  lighter  than  air,  and  sulphur 
dioxide  more  than  twice  as  heavy  (see  Exercise  13). 

Applications:  To  Cases  of  Dissociation.  —  Several  substances 
yield  smaller  values  for  their  densities,  and  therefore  molecular 


APPLICATIONS  OF  MOLECULAR   EQUATIONS  213 

weights,  when  the  densities  are  measured  at  higher  temperatures. 
This  indicates  that  the  molecules  have  become  lighter,  and  can  OD!J 
mean  that  decomposition  has  taken  place  in  consequence  of  the 
heating.  Behavior  of  this  kind  is  shown  both  by  compounds  and  by 
simple  substances. 

For  example,  phosphorus  pentachloride  PC15,  although  a  solid,  can 
be  converted  into  vapor  without  much  .difficulty.  Its  molecular 
weight,  if  it  underwent  no  chemical  change  during  the  volatilization, 
would  be  31  +  177.25  =  208.25.  The  density  actually  observed  at 
300°  gives  by  calculation  not  much  more  than  half  this  value.  The 
direct  inference  from  this  is,  that  the  molecules  have  only  half  the 
(average)  weight  that  we  expected :  or,  in  other  words,  are  twice  as 
numerous  as  we  expected.  The  explanation  is  found  when  we  ex- 
amine the  nature  of  the  vapor  more  closely.  We  find  that  it  is  a  mix- 
ture of  phosphorus  trichloride  and  free  chlorine,  resulting  from  a 
chemical  change  according  to  the  equation  :  PC15  «r±  PClg  +  Cl^  The 
low  value  of  the  density  thus  tells  us  that  dissociation  has  taken  place. 
From  the  value  of  the  density  at  various  temperatures,  we  may  even 
calculate  the  proportion  of  the  whole  material  which  is  dissociated. 
At  300°  it  is  97  per  cent;  at  250°,  80  per  cent;  and  at  200°,  48.5  per 
cent.  Thus,  when  the  temperature  is  lowered,  progressive  recombina- 
tion takes  place  and  the  proportion  dissociated  becomes  less.  Finally 
the  vapor  condenses  and  yields  the  original  solid. 

Again,  sulphur  boils  at  445°,  but  is  easily  vaporized  at  a  tempera- 
ture as  low  as  193°,  under  very  low  pressure.  At  this  temperature  the 
density  of  the  vapor  gives  the  molecular  weight  256  (  =  8  X  32),  and 
the  molecular  formula  S8.  That  is  to  say,  the  G.M.V.  holds  256  g.  of 
the  vapor  at  193°.  At  800°,  however,  the  density  is  only  one-fourth 
as  great,  and  the  G.M.V  holds  only  64  g.  (S2).  This  means  that 
256  g.  now  occupy  four  times  as  large  a  volume  as  before,  and  the  in- 
crease is  additional  to  the  effect  of  the  mere  thermal  expansion,  which  is 
allowed  for  in  the  calculation  and  eliminated.  Hence  the  molecules 
have  dissociated.  At  1700°  the  molecular  formula  is  still  S2,  so  that 
this  shows  the  limit  of  observed  dissociation:  S8  <=±  4S2.  When  the  vapor 
is  cooled,  the  density  increases  once  more  and  at  193°  recovers  com- 
pletely the  greater  value.  Similar  observations  show  that  phosphorus 
vapor  at  313°  is  all  P4,  but  at  1700°  half  of  it  has  dissociated  into  P9. 
Iodine  vapor,  up  to  445°,  is  all  I2.  Beyond  this  temperature  the  den- 
sity diminishes,  and  when  1700°  is  reached  the  vapor  is  all  I,  Thus 
the  molecules  are  diatomic  at  low  temperatures  and  monatomic  at  high 


214  INORGANIC   CHEMISTRY 

ones.  The  densities  of  oxygen,  hydrogen,  and  chlorine  are  not 
measurably  affected  by  heating  to  1700°,  so  that  their  diatomic  mole- 
cules exist  from  temperatures  far  below  0°  up  to  1700°,  and  are  evi- 
dently very  stable. 

Applications:  Finding  the  Atomic  Weight  of  a  Neiv  Ele- 
ment. —  By  way  of  reviewing  the  principles  explained  in  this  chapter, 
let  us  apply  them  to  the  imaginary  case  of  a  newly  discovered  ele- 
ment. The  bromide  of  the  element  is  found  to  be  easy  of  preparation 
and  to  be  volatile.  The  bromide  contains  30  per  cent  of  the  element, 
and  its  vapor  density  referred  to  air  is  11.8.  The  analysis  can  always 
be  made  much  more  accurately  than  the  measurement  of  vapor  density, 
so  that  the  former  number  is  more  trustworthy  than  the  latter. 

To  find  the  equivalent  of  the  element,  that  is,  the  amount  com- 
bined with  80  parts  (the  equivalent)  of  bromine,  we  have  the  propor- 
tion 70  :  80  :  :  30  :  x,  from  which  x  =  34.3.  The  atomic  weight  must 
be  this,  or  some  small  multiple  of  it. 

The  G.M.V.  of  air  weighs  28.955  g.  (p.  191).  Hence  the  same 
volume  of  the  vapor  of  this  bromide,  which  is  11.8  times  as  heavy  as 
air,  will  weigh  28.955  X  11.8,  or  341.67  g.  This  is  therefore  the 
molar  weight  of  the  compound. 

Now  30  per  cent  of  this  is  the  new  element : 

341.67  X  30  -v-  100  =  102.5. 

Three  times  the  equivalent  weight  is  the  multiple  nearest  to  this 
number,  3  X  34.3  =  102.9,  the  difference  being  due  to  error  in  deter- 
mining the  density.  So  long  as  no  other  volatile  compound  is  known, 
we  adopt  this  as  the  atomic  weight.  The  rest  of  the  molar  weight 
(240  parts)  is  bromine.  The  atomic  weight  of  bromine  is  about  80. 
Thus  the  formula  of  the  compound  is  ElBr3,  and  from  this  we  see 
that  the  element  is  trivalent. 

In  case  no  volatile  compound  of  the  element  can  be  formed,  the 
equivalent  is  measured  as  before.  Then  some  of  the  free  simple 
substance  is  made,  say  by  electrolysis,  and  its  specific  heat  is  deter- 
mined. The  sp.  ht.  is  about  0.063.  Application  of  Dulong  and 
Petit's  law  then  gives  the  atomic  weight.  The  product  34.3  X  0.063 
is  equal  to  2.161.  Hence,  the  equivalent  must  be  multiplied  by  3 
to  give  the  atomic  weight,  for  this  raises  the  product  to  6.48,  which 
is  within  the  limits.  Thus  the  value  of  the  atomic  weight  is  102.9. 
as  before. 


APPLICATIONS  OF  MOLECULAR   EQUATIONS  215 

Replies  to  Questions  about  Difficulties.  —  The  beginner  always 
becomes  confused  over  one  or  more  of  the  points  raised  by  the  fol- 
lowing questions : 

Why  was  32  g.  of  oxygen  taken  as  the  standard  for  molecular 
weights,  rather  than  16  g.?  Eead  p.  193  and  footnote  to  p.  197. 

If  02  is  the  smallest  mass  of  oxygen,  why  do  we  have  formulae 
like  H-jO  and  HC10?  O2  is  the  smallest  mass  of  free  oxygen,  but  in 
combination  half  as  much  occurs  in  many  molecules.  Head  pp.  198, 
204,  and  205. 

Why  is  not  the  atomic  weight  of  an  element  ascertained  by  simply 
measuring  the  density  of  the  elementary  substance  ?  Kead  pp.  206 
and  213. 

Can  we  not  deduce  the  valence  of  an  element  from  knowing  the 
number  of  atoms  in  its  molecules,  and  vice  versa?  Some  molecular 
formulae  and  valences  are:  H^1,  02n,  Cy,  Zn11,  also  Hg  (univalent 
and  bivalent),  P4  (trivalent  and  quinquivalent)  and  S8  (bivalent  and 
sexivalent).  There  is  no  relation,  either  observable  or  to  be  expected. 

Do  the  molecular  weights,  oxygen  =  32  and  hydrogen  =  2  mean 
that  the  molecules  of  oxygen  are  larger  than  are  those  of  hydrogen  ? 
This  is  the  ratio  of  their  weights,  but  none  of  the  phenomena  dis- 
cussed in  this  chapter  are  influenced  appreciably  by  their  relative 
sizes,  and  therefore  none  of  them  give  any  information  on  the  subject. 
Eead  the  footnote  to  p.  195. 

Exercises.  —  1.  The  weight  of  1  1.  of  gas  at  0°  and  760  mm.  is 
5.236  g.  What  is  the  density  referred  to  air  and  to  hydrogen,  and 
what  is  the  molecular  weight  (pp.  191,  194)  ? 

2.  The   density  of  a  gas,  referred  to  air,  is  6.7.      What   is   the 
weight  of  1  1.  (p.  191),  and  what  is  the  molecular  weight  (p.  214)? 

3.  The  molecular  weight  of  a  substance  is  65.     What  is  the  den- 
sity referred  to  air,  and  what  is  the  weight  of  1  1.  ? 

4.  The  chloride  of  a  new*  element  contains  38.11  per  cent  of  chlo- 
rine and  61.89  per  cent  of  the  element.     The  vapor  density  of   the 
compound  referred  to  air  is  12.85.     What  is  the  atomic  weight  of  the 
element,  so  far  as  investigation  of  this  one  substance  can  give  it  (p.  196)? 
What  is  its  valence  ? 

5.  If  the  molecular  weight  of  oxygen  were  taken  as    100,  what 
would  be  the  volume  of  the  G.M.V.  (p.   191).     On   the  same    scale 
what  would  be  the  molecular  weight  of  water,  and  what  would  be  the 
atomic  weights  of  hydrogen  and  chlorine  (pp.  191,  198)  ? 


216  INORGANIC    CHEMISTRY 

6.  In  future  nothing  but  molecular  formulae  of  free  elements  must 
be  used  (p.  210).     Write  in  molecular  form  ten  of  the  equations  in- 
volving gases  which  are  found  in  the  preceding  chapters. 

7.  If  a  new  form  of  oxygen  were  found,  such  that  one  volume  of 
it  required  four  volumes  of  hydrogen  to  produce  water,  what  would 
be  its  molecular  formula  (p.  207)?     What  would  be  the   weight  of 
22.4  1.  ? 

8.  What  volume  of  oxygen  at  10°  and  750  mm.  is  obtainable  by 
heating  50  g.  of  barium  peroxide  (pp.  64,  209,  210)  ? 

9.  What  volume  of  oxygen  at  20°  and  760  mm.  is  required  to  con- 
vert 16  g.  of  iron  into  dehydrated  rust  (Fe203)  (p.  210)  ? 

10.  Write  out   the   molecular  equations  for   the   interactions   of 
methane  and  chlorine  (pp.  176,  209) ;  and  for  the  burning  of  phosphorus 
(vapor)  in  oxygen  (pp.  198,  204,  211).     Deduce  the  volume  relations  of 
the  initial  substances,  and  of  the  products,  at  various  temperatures  in 
each  case. 

11.  Write  out  the  molecular   equations    for   the   interactions   of 
acetylene  and  oxygen  (p.  212),  and  of  alcohol  vapor  (b.-p.  78°)  and 
oxygen.      Deduce  the  volume  relations  of  the  initial  substances  and 
of  the  products  at  0°  and  at  100°  in  each  case. 

12.  The  molecular  weight  of  cyanogen  is  52.08.     What  is  its  den- 
sity referred  to  air,  and  what  the  weight  of  1  1.  at  0°  and  760  mm.?  It 
contains  46.08  per  cent  carbon  and  53.92  per  cent  nitrogen.     What  is 
the  formula  of  the  substance  (p.  55)  ?     Exploded  with  oxygen  it  forms 
carbon  dioxide  and  free  nitrogen.     What  will  be  the  relative  volumes 
of  the  materials  before  and  after  the  interaction  (p.  212)? 

13.  What  are  the  relative  weights  of  equal  volumes  of  hydrogen 
sulphide  (H2S),  and  hydrogen  iodide  (HI),  compared  with  air  (p.  212)? 

14.  At  1700°  the  average  molecular  weight  of  phosphorus  is  91 
(p.  213).     What  percentage  of  molecules  of  P4  has  been  dissociated 
into  P2  ? 

15.  Show  that,  if  an  element  has  more  than  one  equivalent  weight, 
the  atomic  weight  must  be  some  multiple  of  each  of  the  equivalents 
by  a  whole  number. 

16.  Prove,  without  the  use  of  anything  hypothetical,  that  16  is 
preferable  to  8  for  the  atomic  weight  of  oxygen,  because  the  smaller 
number  involves  a  fractional  value  for  the  atomic  weight  of  hydrogen. 


CHAPTER   XIII 
THE   ATOMIC   HYPOTHESIS 

To  determine  the  nature  of  chemical  phenomena  requires,  as  we 
have  found,  very  elaborate  experimentation.  And  this  has  to  be 
followed  by  still  more  elaborate  reasoning  before  a  systematic  state- 
ment of  the  precise  nature  of  the  change  can  be  made.  Yet,  when 
all  this  has  been  done,  we  are  still  unable  to  form  a  clear  conception 
of  what  manner  of  procedure  the  change  follows,  for  the  details  are 
entirely  inaccessible  to  observation.  We  should  like  to  know  pre- 
cisely how  chemical  union  is  consummated,  and  how  chemical  exchange 
is  carried  out.  We  should  like  to  account  for  the  fact  that  the  auto- 
matically adjusted  proportions  of  the  materials  used  in  every  chemical 
change  are  entirely  uninfluenced  by  temperature  and  other  conditions. 
Above  all,  we  should  like  to  know,  if  possible,  what  state  of  affairs 
determines  the  employment  of  an  individual  unit  weight  by  each 
element  in  all  its  combinations.  None  of  these  questions  can  be 
answered,  however,  because  nothing  can  be  seen  which  suggests  any 
answer. 

As  usual  in  cases  of  this  kind  we  construct  an  imaginary  mechanism, 
a  formulative  hypothesis  (p.  141)  to  account  for  the  facts.  In  doing 
so,  we  are  not  under  the  illusion  that  we  are  discovering  the  actual 
machinery.  We  realize  that  we  are  simply  making  a  sort  of  diagram 
which  will  assist  our  thought  about  the  thing  itself.  Now  a  slight 
addition  to  the  molecular  hypothesis  readily  furnishes  precisely  what 
we  need. 

Atomic  Hypothesis.  —  According  to  the  molecular  hypothesis,  all 
matter  is  made  up  of  small  discrete  particles,  each  of  which  has  the 
same  composition  as  has  the  body  as  a  whole.  It  is  difficult,  there- 
fore, to  avoid  the  conception  that  the  different  materials  in  each  com- 
pound molecule  are  more  or  less  distinct  entities  also.  Hence  we  make 
this  the  basis  of  a  new  hypothesis,  and  attribute  to  these  constituent 
parts  of  molecules  the  properties  of  the  chemical  unit  weights  which 
are  closely  related  to  them.  These  parts  must  move  from  one  state  of 

217 


218  INORGANIC   CHEMISTRY 

combination  to  another  without  alteration  in  their  mass.  Since  they 
may  be  restored  to  their  free  condition  and  recombined  as  often  as  we 
choose,  without  impairment  of  their  individuality,  each  kind  must  be 
composed  of  a  distinct  variety  of  matter.  These  two  are  almost  the 
only  qualities  which  the  facts  thus  far  presented  justify  us  in  attri- 
buting to  them.  We  must,  therefore,  carefully  avoid,  for  the  present, 
the  introduction  of  unnecessary  complications  by  inventing  for  them 
any  other  properties  gratuitously. 

These  parts  of  molecules  which  we  thus  suppose  to  be  permanent, 
coherent  masses,  are  named  atoms.  This  word  signifies  objects  which 
are  not  disintegrated  (Gk.  arofios,  uncut,  i.e.,  not  yet  cut).  The  relative 
weights  of  these  imaginary  masses  are  numerically  the  same  as  the 
chemical  unit  weights  (p.  50),  and  hence,  in  terms  of  this  hypothesis, 
we  call  the  latter  atomic  'weights. 

It  should  be  noted  that  although  atoms,  like  molecules,  are  fictions, 
the  atomic  weights,  since  they  are  measured  by  experimental  methods, 
are  real.  They  originally  received  this  name  when  Dalton,  an  English 
schoolmaster  of  Manchester,  succeeded  in  unraveling  the  complications 
of  the  chemical  composition  of  substances  by  the  help  of  this  very 
hypothesis,  and  realized  for  the  first  time  (1805)  that  the  possession 
by  all  elements  of  individual  chemical  unit  weights  lay  at  the  basis  of 
the  whole  system. 

We  may  sum  up  all  that  the  facts  require  us  to  assume  about  atoms 
by  saying :  Atoms  are  the  units  of  which  molecules  are  aggregates. 
Those  of  like  kind  have  equal  masses,  and  differ  from  those  of  other 
kinds  both  in  mass  and  the  kind  of  material  of  which  they  are  made. 
The  fundamentally  different  kinds  of  materials  are  the  chemical 
elements.  These  statements  include  the  whole  atomic  hypothesis. 

It  is  to  be  noted,  particularly,  that  there  are  no  facts  in  chemistry  which  prove 
that  atoms  are  incapable  of  disintegration.  All  we  know  is  that  in  ordinary  chemi- 
cal changes  they  are  transferred  as  wholes.  It  would  only  occasion  the  risk  of 
conflict  with  facts  still  to  be  discovered  if  we  elaborated  our  fiction  any  further 
than  is  absolutely  necessary.  Indeed,  the  phenomena  of  radio-activity  have  already 
compelled  us  to  suppose  that  there  are  particles  much  smaller  than  atoms  (see  below). 
Without  this  assumption  the  new  facts  cannot  be  accounted  for  in  harmony  with 
the  molecular  hypothesis. 

The  idea  that  matter  is  composed  of  small  particles  is  a  very  ancient  one. 
Not  even  Dalton,  however,  although  he  used  this  conception  continually  as  a 
means  of  thought  about  chemical  and  physical  phenomena,  made  any  distinction 
between  atoms  and  molecules.  This  distinction  was  introduced  later  by  Avogadro 
(1811).  Yet,  without  this  refinement,  continual  thought  of  the  behavior  of  matter 
as  consisting  in  transactions  between  small  particles  led  Dalton  to  see  that  it  was 


THE   ATOMIC    HYPOTHESIS  219 

probable  that  individual  unit  weights  for  each  element  must  exist.  The  discovery 
that  they  did  exist  soon  followed.  The  numbers  which  Dalton  actually  gave  out, 
aside  from  the  considerable  experimental  inaccuracies  attached  to  them,  were  often 
equivalents  and  not  modern  atomic  weights.  It  was  after  the  publication  of 
Dalton's  ideas  that  Gay-Lussac  discovered  the  law  of  combining  volumes  (1808), 
and  until  this  law  was  discovered  there  was  no  criterion  for  fixing  the  values  of  the 
atomic  weights.  Gay-Lussac,  at  the  end  of  his  paper,  pointed  out  that  his  dis- 
covery formed  an  important  confirmation  of  Dalton's  views.  Strange  to  say, 
Dalton  himself  refused  to  accept  Gay-Lussac's  law,  and  so  rejected  the  very  means 
by  which  his  own  principle  of  chemical  unit  weights  came  eventually  to  be  acknowl- 
edged as  one  of  the  foundation  stones  of  the  science.  On  the  other  hand,  Dalton's 
fellow  countrymen  and  contemporaries  accepted  the  principle  of  unit  weights, 
but  rejected  the  atomic  hypothesis  by  the  help  of  which  Dalton  had  reached  them  ! 
Thus  Sir  Humphry  Davy  called  them  "  proportions  "  instead  of  atomic  weights, 
and  Wollaston  preferred  the  word  "  equivalents." 

The  Atomic  Hypothesis  and  Chemical  Change.  —  The  change 
from  chemical  units  to  atoms  is  so  slight  that  the  application  of  this 
hypothesis  to  the  description  of  chemical  phenomena  is  very  readily 
made.  The  description  gains  in  concreteness,  however,  from  the 
change.  Thus,  we  consider  iron  from  every  source  to  be  made  of 
minute  portions  of  iron  matter  which  are  all  alike  in  weight,  and  pre- 
sumably also  in  their  other  properties.  Similarly,  all  specimens  of 
sulphur  are  made  of  minute  particles  of  sulphur  having  exactly  the 
same  weight.  The  weight  of  a  sulphur  atom,  however,  is  different 
from  that  of  an  iron  atom.  When  visible  portions  of  the  two  sub- 
stances unite,  we  conceive  the  operation  to  consist  in  the  union  of 
each  atom  of  iron  with  one  atom  of  sulphur  to  produce  a  molecule  of 
ferrous  sulphide.  The  consummation  of  this  union  between  multi- 
tudes of  pairs  of  the  respective  kinds  of  atoms  in  every  second  of  time 
results  in  a  chemical  transformation  whose  progress  is  perceptible  by 
the  senses. 

In  more  complex  chemical  changes,  a  further  correspondence  be- 
tween the  hypothesis  and  the  facts  comes  to  our  notice.  For  example, 
when  sulphuric  acid  acts  upon  sodium  chloride  to  produce  hydrogen 
chloride : 

NaCl  +  H2S04  ->  HC1  +  NaHS04, 

one  part  by  weight  of  hydrogen  takes  the  place  of  23  parts  by  weight 
of  sodium,  and  combines  with  35.45  parts  of  chlorine  to  form  the 
hydrogen  chloride.  This  is  the  way  we  state  the  change  when  we 
refer  to  measurable  quantities.  According  to  this  hypothesis,  the 
operation  consists  in  the  repetition,  millions  of  times  over  within  a 


220  INORGANIC   CHEMISTRY 

small  amount  of  material,  of  the  substitution  of  one  atom  of  hydrogen 
for  one  atom  of  sodium  to  form  a  molecule  of  hydrogen  chloride. 
The  special  fact  which  we  notice  is  that  the  atom  of  hydrogen  suffices 
exactly  to  occupy  the  place  of  the  atom  of  sodium.  If  it  were  too 
large,  then  a  portion  of  each  atom  of  hydrogen  would  remain  unused, 
and  so  some  free  hydrogen  would  be  produced  along  with  the  other 
products.  If  it  were  too  small,  then  some  atoms  of  hydrogen  would 
be  consumed  in  making  up  the  material  of  the  others  to  the  right 
amount,  and  hence  some  chlorine  might  fail  to  receive  any  hydrogen 
in  combination  and  escape  in  a  free  condition.  Neither  of  these  things 
is  observed  to  take  place,  however.  The  remarkable  fact  about  this, 
and  all  other  double  decompositions  of  the  same  sort,  is  that  the  little 
masses  of  the  various  elements  exchange  places  without  any  altera- 
tions to  fit  the  new  compound  being  required.  Hence  our  assumption 
that  the  atoms  are  permanent,  coherent  wholes.  This,  of  course,  is 
simply  stating  in  terms  of  the  atomic  hypothesis  the  facts  which 
underlie  the  conception  of  unit  weights.  A  chemical  phenomenon 
as  we  observe  it,  then,  is  imagined  to  consist  in  some  systematic  libera- 
tion, combination,  or  exchange  of  atoms,  according  to  a  definite  scheme, 
and  repeated  many  millions  of  times  (with  every  molecule)  in  a  body  or 
mixture  of  bodies. 

The  Atomic  Hypothesis  and  the  Quantitative  Laws.  —  The 

idea  of  atoms  simply  crystallizes  somewhat  more  definitely  the  con- 
ception of  chemical  unit  weights.  Hence,  it  follows  of  necessity  that 
the  quantitative  laws  of  chemical  combination,  out  of  which  the  latter 
arose,  will  be  found  to  be  entirely  in  harmony  with  the  atomic  hypothe- 
sis. The  definite  composition  of  each  compound,  for  example,  corre- 
sponds to  the  hypothesis  that  each  substance  is  made  up  of  a  specific 
kind  of  molecules,  all  of  which,  in  turn,  contain  the  same  kind  and 
number  of  atoms. 

The  conception  of  valence  (p.  101)  suggests,  in  terms  of  this 
hypothesis,  that  some  atoms  unite  with  but  one  other  atom,  habitually 
(NaCl).  Some,  however,  unite  with  two  of  the  first  kind  (ZnCl2),  or 
with  one  other  of  their  own  kind  (ZnO),  still  others  with  three  atoms 
of  the  first  kind  (A1C13),  and  so  forth.  In  other  words,  it  involves 
the  assumption  that  each  kind  of  atom  has  a  limited  capacity  for 
holding  other  atoms  in  combination.  Thus,  taking  the  most  crudely 
mechanical  view  of  the  matter,  we  might  elaborate  the  hypothesis  by 
suggesting  that  there  is  a  limit  to  the  number  of  points  at  which 


THE    ATOMIC   HYPOTHESIS  221 

atoms  may  be  attached  to  one  another.  When  one  atom  of  chlorine 
is  attached  to  one  of  sodium,  the  combining  capacity  of  each  is  ex- 
hausted. When  one  atom  of  hydrogen  is  attached  to  one  atom 
of  oxygen,  one  combining  capacity  still  remains  (H — 0 — ),  and 
can  be  satisfied  by  one  more  atom  of  hydrogen  or  of  some  other 
element. 

The  fact  that  the  vapor  density  of  the  elements  is  often  greater 
than  that  of  the  compounds  in  which  they  are  contained,  a  fact  which 
led  us  (p.  205)  to  the  conception  that  there  are  two  or  more  unit 
weights  in  one  molecule,  becomes  more  concrete  in  view  of  this  hy- 
pothesis. It  is  as  easy  to  understand  that  there  may  be  two  or  more 
atoms  of  the  same  kind,  as  two  or  more  of  different  kinds,  united  in 
one  molecule. 

Thus  the  general  scheme  according  to  which  chemical  changes 
occur,  as  well  as  all  the  quantitative  relations  in  chemical  phenomena, 
are  admirably  described  when  we  think  of  those  phenomena  as  con- 
stituted by  transferences  of  atoms  possessing  definite  weights. 

Equations  and  the  Atomic  Hypothesis.  —  Since  equations  are 
simply  records  of  the  kinds  and  quantities  of  matter  taking  part  in 
chemical  changes,  they  require  no  addition  to  the  atomic  hypothesis. 
The  symbols,  which  originally  represented  unit  weights  of  the  various 
elements,  may  stand  equally  well  for  atoms.  Hence,  in  the  current 
language  of  chemistry,  potassium  chlorate  (KC108)  is  composed  of 
one  atom  each  of  potassium  and  of  chlorine  and  three  atoms  of  oxy- 
gen in  each  molecule.  The  word  "  atom  "  stands  for  unit  weight,  and 
the  word  "  molecule  "  for  molecular  weight. 

Concretely,  when  these  terms  are  used,  the  chemical  equation 
represents  one  minute  specimen  of  the  change  which  is  taking  place 
throughout  the  whole  mass.  It  shows  enough  molecules  and  atoms  to 
furnish  a  complete  example  of  what,  repeated  millions  of  times,  will 
constitute  a  chemical  change  in  a  visible  amount  of  material.  There 
are  a  few  details  which  this  point  of  view  suggests.  For  example, 
the  equation  HgO  — >  Hg  +  0  shows  the  weights,  but  does  not  include 
a  complete  sample  of  the  materials  in  the  point  of  view  of  this 
hypothesis.  The  smallest  fragment  of  free  oxygen  which  we  can 
obtain  is  made  of  two  atoms.  Hence  two  molecules  of  mercuric  oxide 
are  required  to  furnish  them.  Thus  the  minimum  number  of  mole- 
cules which  would  suffice  for  carrying  out  this  change  on  a  molecular 
scale  is  2HgO  — >  2Hg  -f-  02.  Similarly,  since  the  smallest  discrete 


222  INORGANIC   CHEMISTRY 

portions  of  free  hydrogen  and  chlorine   contain  two  atoms  each,  we 
must  write  the  equation  H2  +  CLj  ->  2HC1  (cf.  p.  207). 

May  other  Properties,  Besides  Definite  Mass  and  Specific 
Material,  be  attributed  to  an  Atom  ?  —  When  this  question  is  raised, 
the  ground  becomes  less  certain.  A  few  examples  will  show  the  dif- 
ficulties encountered  and  how  far  useful  suggestions  have  been  made. 

Since  mercury  and  iodine  in  a  gaseous  condition  are  composed  of  single  atoms 
(p.  205),  all  atoms  must  be  perfectly  elastic  just  as  molecules  were  assumed  to  be 
(p.  129).  But  this  assumption  is  unnecessary  in  chemistry,  for  it  throws  no  light 
upon  chemical  behavior. 

The  internal  constitution  of  atoms  remained  for  a  century  rather 
indefinite.  As  fictions,  whose  properties  depended  entirely  upon 
those  of  the  chemical  units  which  they  represented,  they  were  not 
regarded  as  ever  undergoing  any  change  in  the  amount  of  material 
they  contained.  Recently,  however,  the  peculiar  radiations,  which 
proceed  from  certain  minerals  containing  uranium  and  thorium,  have 
been  discovered  to  be  composed  of  minute  particles  (see  Radium). 
It  has  been  found  possible,  even,  to  estimate  the  size  which  these 
particles  must  possess  as  compared  with  atoms,  in  order  that  they  may 
exhibit  the  properties  which  are  peculiar  to  them.  Different  methods 
give  somewhat  different  values.  But  it  appears  that  some  of  these 
particles,  which  are  called  corpuscles  (sometimes  electrons)  to  dis- 
tinguish them  from  atoms,  are  of  the  same  size  from  whatever  element 
they  are  formed,  and  that  their  mass  is  about  one-thousandth  of  that  of 
a  hydrogen  atom.  The  relation  is  of  this  order,  at  all  events.  Their 
behavior  also  compels  us  to  suppose  that,  even  in  the  atom  from  which 
they  issue,  they  are  scattered  at  wide  distances  from  one  another.  Thus, 
Professor  J.  J.  Thomson  suggests  that,  if  the  hydrogen  atom  were  as 
large  as  a  church,  the  thousand  or  so  electrons  which  it  contains 
might  be  compared  to  that  number  of  full  stops  (.)  scattered  through- 
out the  whole  edifice.  We  have  thus  been  compelled  by  the  facts  to 
add  to  the  molecular  hypothesis,  first  an  atomic  hypothesis,  and  then 
a  corpuscular  hypothesis.  Our  atoms  turn  out  to  be  clusters  of 
particles  after  all  (p.  220).  But  the  atom  (chemical  unit)  of  ordinary 
chemical  change  is  not  thereby  affected,  however  many  corpuscles 
may  be  thrust  into  it  by  other  considerations. 

In  one  direction,  namely,  that  of  accounting  for  the  properties  of 
compounds,  no  attempt  has  been  made  to  adapt  the  atomic  hypothesis 
so  as  to  explain  the  facts.  Thus,  two  hydrogen  atoms  in  a  molecule 


THE   ATOMIC    HYPOTHESIS  223 

give  a  gas  almost  insoluble  in  water  ;  two  chlorine  atoms,  a  gas  which 
is  moderately  soluble ;  but  one  of  each  gives  hydrogen  chloride,  which 
dissolves  in  water  in  extraordinary  quantities.  So  also,  colorless 
substances  give,  by  chemical  union,  strongly  colored  ones,  and  odorless 
substances,  by  chemical  union,  strongly  odorous  ones.  Putting  pieces 
of  iron  and  sulphur  side  by  side  causes  absolutely  no  change  in  the 
properties  of  either.  And  yet  this  hypothesis  compels  us  to  assume 
that  if  the  particles  are  made  fine  enough,  and  placed  close  enough  to 
one  another,  the  individual  properties  of  the  constituents  will  entirely 
disappear.  Hitherto  we  have  failed  to  think  of  any  qualities  which 
might  be  attributed  to  the  atoms  in  order  to  account  for  facts  of  this 
class.  Why  should  oxygen  (02)  and  ozone  (08)  be  so  different  in 
behavior  although  the  atomic  theory  hints  at  nothing  but  a  substitu- 
tion of  three  atoms  for  two  ?  What  atomic  properties  shall  account 
for  the  difference  between  red  and  yellow  phosphorus  ? 

Usually,  in  explaining  these  matters,  we  forsake  atoms  and  employ 
the  conception  of  energy.  We  say  that  ozone  has  more  energy  than 
oxygen,  and  yellow  phosphorus  than  red  phosphorus.  Here  the  con- 
ception of  atoms  as  a  means  of  accounting  for  properties  is  subordi- 
nate to  that  of  energy.  A  tendency  is  plainly  perceptible,  indeed,  not 
merely  to  supplement  the  atomic  theory  by  the  notion  of  energy,  but 
to  allow  energy  to  supplant  it  entirely,  and  that  not  only  in  this  but 
in  other  connections. 

There  are  certain  special  peculiarities  in  the  relation  between  com- 
position and  properties,  however,  to  the  explanation  of  which  the 
atomic  theory  has  been  successfully  adapted.  Thus,  certain  com- 
pounds precisely  similar  in  composition,  such  as  barium  peroxide 
(Ba02)  and  lead  dioxide  (Pb02),  show  entirely  different  chemical  be* 
havior.  The  former  of  these  when  treated  with  an  acid  gives  hydrogen 
peroxide  (q.v.\  while  the  latter  gives  only  water  and  oxygen.  If  it  is 
merely  a  question  of  composition,  the  substances  should  behave  alike, 
So,  also,  we  often  have  two  or  more  compounds  identical  in  nioleculaJ 
weight,  and  in  the  elements  and  numbers  of  units  which  they  contain, 
which  are  nevertheless  totally  different  in  physical  and  chemical  prop* 
erties.  To  account  for  this  difference  we  have  found  it  convenient 
to  suppose  that,  although  the  atoms  contained  in  the  two  or  more 
kinds  of  molecules  are  of  the  same  numbers  and  kinds,  they  are  iff 
different  geometrical  arrangement  towards  one  another  (see  Urea). 
Sometimes  one  of  these  substances  can  be  made  directly  from  the  other, 
and  this  gives  us  that  variety  of  chemical  change  which  was  named 


224  INORGANIC   CHEMISTRY 

"  internal  rearrangement  "  (p.  15).  In  discussing  facts  like  the  above 
in  terms  of  the  atomic  hypothesis  we  speak  of  the  structure  of  the 
molecule  and  the  constitution  of  the  substance,  and,  to  represent  our 
conclusions,  we  use  graphic  formulae  (see  Constitutions  of  oxygen 
acids  of  chlorine,  hydrogen  peroxide,  sulphuric  acid,  formic  acid  and 
urea). 

It  is  becoming  increasingly  difficult  to  maintain  the  distinction  which  formerly 
seemed  very  clear  between  physical  and  chemical  phenomena.  The  atomic  hy- 
pothesis serves  the  purpose  in  a  rough  way  for  the  simpler  cases.  If  the  atoms 
are  undisturbed,  and  therefore  the  chemical  properties  remain  unchanged,  the 
substance  has  undergone  a  physical  change  only.  If,  however,  the  atoms  are 
altered  in  their  relations  to  one  another,  either  by  rearrangement,  or  in  some  more 
drastic  manner,  the  change  is  a  chemical  one.  Now,  however,  so  many  shades  of 
difference  among  the  examples  that  come  under  the  head  of  internal  rearrange- 
ment have  been  found,  that  the  atomic  hypothesis  is  applied  with  constantly 
increasing  difficulty  (cf.  p.  16).  Whether  it  will  be  possible  to  apply  this 
hypothesis  consistently  in  this  region  remains  yet  to  be  seen. 

Summing  this  up,  we  find  that  all  the  suppositions  the  chemist 
makes  are  :  That  an  atom  has  a  specific  mass  and  consists  of  a  specific 
kind  of  material,  and,  sometimes,  that  the  atoms  in  a  molecule  are 
arranged  with  reference  to  one  another  in  space  in  some  definite  way. 
After  making  this  fornmlative  hypothesis,  however,  he  knows  no  more 
than  he  did  before  about  the  real  mechanism  of  chemical  change. 

Is  the  Atomic  Hypothesis  a  Fact.  — The  language  of  chemists  has 
become  so  saturated  with  the  phraseology  of  the  atomic  and  molecular  hypotheses, 
that  we  speak  in  terms  of  atoms  and  molecules  as  if  they  were  objects  of  immedi- 
ate observation.  It  must  be  reiterated,  therefore,  that  this  language  is  figurative, 
and  must  not  be  taken  literally.  The  atomic  hypothesis  provides  a  convenient  form 
of  speech,  which  successfully  describes  many  of  the  facts  in  a  metaphorical  manner. 
But  the  handy  way  in  which  the  atomic  hypothesis  lends  itself  to  the  representa- 
tion of  the  characteristic  features  of  a  chemical  change  falls  far  short  of  constituting 
a  proof  that  atoms  have  any  real  existence.  The  restatements  which  we  have 
made  of  some  of  the  characteristics  of  chemical  change  in  this  chapter  depend  upon 
the  continued  reign  of  the  atomic  hypothesis.  The  definitions  of  the  same  things 
given  previously  were  based  upon  experiment  and  are  of  permanent  value,  so  that 
no  question  can  be  allowed  to  arise  as  to  which  of  the  two  definitions  in  each  case 
is  the  more  important  one. 

The  existence  of  the  unit  weights  (atomic  weights)  will  probably  appear  to  come 
nearer  to  furnishing  verification  of  the  hypothesis  that  matter  is  composed  of 
atoms  than  any  other  fact.  It  may  seem  unthinkable  that  each  element  should  use 
the  same  proportions  by  weight  in  entering  into  many  different  combinations  if  it  is 
not  actually  done  up  in  packets  whose  weight  is  represented  by  this  number.  Such 
an  inference,  however,  is  not  really  justified.  An  illustration  will  make  this  clear. 
If  we  were  unable  to  see  wheat  close  at  hand,  so  as  to  distinguish  its  structure,  and 


THE  ATOMIC   HYPOTHESIS  225 

our  information  was  confined  to  observation  from  a  distance  and  the  reading  of 
market  quotations,  we  should  be  led  to  infer  from  the  latter  that  it  was  a  substance 
which  was  always  done  up  in  bushels.  Yet  we  know  that  this  inference  would  be 
entirely  incorrect.  The  substance  really  exists  in  the  form  of  bushels  at  the 
moment  of  measurement  only.  So  there  might  be  imagined  properties,  at  present 
unknown  to  us,  which  directed  the  quantitative  selection  of  material  for  chemical 
change  and  rejected  the  excess,  without  the  existence  of  any  permanent  segregra- 
tion  into  pieces  of  unalterable  dimensions.  The  only  bushels  and  ounces  which 
we  have  are  in  the  measuring  apparatus  and  not  in  the  material  measured  ;  so  the 
only  atomic  weights  may  be  in  the  properties  controlling  chemical  combination, 
and  not  in  the  matter  combining. 

Exercises.  —  1.  In  previous  chapters  our  definitions  have  been  ex- 
perimental. In  imitation  of  the  definitions  of  the  law  of  definite  pro- 
portions and  of  valence  (p.  220),  give  theoretical  definitions  of  the 
following,  in  terms  of  the  atomic  hypothesis :  Physical  and  chemical 
phenomenon,  multiple  proportions,  chemical  unit  weight,  molecular 
weight,  element,  compound,  symbol,  formula,  equation. 

2.  Criticize  the  definitions  :  The  atomic  weight  of  an  element  is  the 
smallest  portion  of  that  element  which  takes  part  in  chemical  change. 
An  atom  is  the  smallest  particle  that  can  be  conceived. 

3.  Define  all  the  varieties  of  chemical  change  (p.  187)  in  terms  of 
the  atomic  hypothesis. 


CHAPTER  XIV 

THE   HALOGEN  FAMILY 

THE  elements  to  which  we  have  so  far  devoted  most  attention  have 
been  oxygen,  hydrogen,  and  chlorine.  If  we  recall  the  chemical  prop- 
erties and  relations  of  these  elements  we  shall  recognize  the  fact  that 
they  all  possess  very  distinct  individualities. 

The  Chemical  Relations  of  Elements.  —  Hydrogen  is  the  sub- 
stance (p.  109)  which  unites  readily  with  oxygen  and  chlorine,  less 
readily  with  other  non-metals,  and  scarcely  at  all  with  rnetals.  Oxy- 
gen and  chlorine  resemble  one  another  somewhat  in  the  greatness  of 
their  chemical  activity  and  the  variety  of  free  elements  with  which 
they  are  capable  of  uniting,  but  differ  markedly  in  what  we  have 
called  their  chemical  relations  (p.  177).  The  resulting  compounds 
belong,  in  fact,  to  quite  different  classes  —  oxygen  forms  oxides,  chlo- 
rine forms  chlorides  —  and  elements  are  considered  similar  only  when 
they  resemble  one  another  in  chemical  relations,  and  produce,  by  com- 
bination -with  the  same  element,  compounds  having  similar  chemical 
properties.  Thus  the  common  oxide  of  hydrogen,  water,  is  a  neutral 
substance,  and  is  chemically  rather  indifferent.  The  chloride  of  hydro- 
gen in  aqueous  solution  is  a  strong  acid  and  is  chemically  very  active.* 
If  all  the  other  chemical  elements  differed  from  one  another  as  much  as 
do  these  three,  the  study  of  the  chemical  elements  would  be  tedious  and 
tiresome,  since  we  should  be  denied  the  satisfaction  of  tracing  resem- 
blances, and  the  elements  would  be  incapable  of  classification.  In 
reality,  however,  we  find  that  they  are  not  incapable  of  being  grouped 
together  in  sets.  They  are  classified  according  to  the  kind  of  sub- 
stances with  which  they  combine  and  the  chemical  nature  of  the  prod- 
ucts. In  some  families  the  resemblance  is  close,  in  others  less  close. 

*  The  difference  between  oxides  and  chlorides  is  seen  in  their  behavior,  and  will 
show  itself  as  we  proceed.  We  have  already  learned,  however,  that  oxides  often 
unite  with  water  to  form  acids  or  bases  (p.  119).  Chlorides  do  not  unite  with 
water  to  form  new  substances  with  marked  characteristics  (c/.  p.  121).  They 
belong  to  the  large  class  of  compounds  designated  salts  (g.u.)- 


THE    HALOGEN   FAMILY  227 

The  present  group  is  of  the  former  class,  and  will  serve,  therefore,  as 
a  convenient  beginning  in  the  work  of  tracing  relations  between  the 
elements  and  in  classifying  the  facts  of  descriptive  chemistry. 

The  Chemical  Relations  of  the  Halogens.—  The  bromide  (NaBr), 
iodide  (Nal),  and,  to  a  less  extent,  the  fluoride  (NaF),  of  sodium, 
resemble  sodium  chloride  (NaCl)  in  appearance  and  behavior.  From 
the  fact  that  chlorine,  bromine,  iodine,  and  fluorine  are  thus  all  able  to 
produce  substances  which  in  composition  and  chemical  actions  resem- 
ble salt,  they  are  known  as  the  halogens  (Gkc  oAs,  salt  ;  yewav,  to  pro- 
duce), and  their  compounds  are  named  the  halides.  The  halogens,  as 
the  above  formulge  show,  are  univalent.  They  all  form  compounds 
with  hydrogen,  and  these  compounds  closely  resemble  hydrogen  chlo- 
ride (q.v.).  For  example,  they  are  colorless,  they  are  gases  (with  the 
exception  of  hydrogen  fluoride  which  is  a  very  volatile  liquid),  they 
are  very  soluble  in  water,  and  their  solutions  are  acids.  Other  rela- 
tions will  be  given  in  a  summary  at  the  end  of  the  chapter. 

BROMINE. 

Occurrence.  —  The  compounds  of  chlorine,  bromine,  and  iodine  usu- 
ally occur  together  in  nature,  while  the  compounds  of  fluorine  are  not 
found  in  the  same  sources.  Bromine  occurs  chiefly  in  the  form  of  the 
bromides  of  sodium  and  magnesium,  in  the  upper  layers  of  the  natural 
beds  of  rock  salt. 

Preparation.  —  In  the  chemical  point  of  view  there  are  three  dis- 
tinct ways  in  which  bromine  is  made.  The  first  of  these  is  closely 
related  to  the  common  method  of  preparing  chlorine  (p.  172).  As 
hydrobromic  acid,  unlike  hydrochloric  acid,  is  not  formed  extensively 
in  connection  with  any  chemical  industry,  potassium  bromide,  the 
most  accessible  compound  of  bromine,  is  treated  with  concentrated 
sulphuric  acid,  and  the  product  is  oxidized  with  powdered  manganese 
dioxide  in  one  operation.  The  whole  action  may  be  represented  in  a 
single  equation  (see  next  section),  as  follows  : 


2KBr  -f  3H2S04  +  Mn02  ->  MnS04  +  2KHS04  +  2KP  +  Br 


2. 


Bromine  being  a  volatile  liquid,  while  the  two  sulphates  are  involatile, 
its  vapor  passes  off  along  with  a  little  water  when  the  above  mixture 
is  heated.  It  is  condensed  in  a  worm-tube  surrounded  by  cold  water. 


228  INORGANIC   CHEMISTRY 

The  second  method  of  preparing  bromine  depends  on  the  fact  that 
chlorine  is  a  more  active  element  and  displaces  bromine  from  combi- 
nation. When,  therefore,  chlorine  is  passed  into  a  solution  of  potas- 
sium or  sodium  bromide,  potassium  or  sodium  chloride  is  formed  and 
the  bromine  liberated  : 

2NaBr  +  CL,  ->  2NaCl  +  Br2. 

When  the  liquid  is  warmed,  the  bromine  passes  off  along  with  a  part 
of  the  water,  and  may  be  condensed  as  before. 

Aqueous  solutions  of  soluble  bromides  may  be  decomposed  by 
means  of  a  current  of  electricity.  The  bromine  is  set  free  at  the 
positive  electrode. 

The  whole  of  the  bromine  used  in  commerce  is  at  present  manu- 
factured in  the  first  two  of  these  ways.  Two-thirds  of  the  supply  is 
obtained  from  Stassfurt,  where,  after  the  extraction  of  the  potas- 
sium chloride  from  the  impure  carnallite  (KC1,  MgCl2,  6H20),  the 
mother-liquid  is  found  to  contain  the  more  soluble  sodium  and  mag- 
nesium bromides  in  considerable  quantities.  The  warm  mother-liquor 
trickles  down  over  round  stones  in  a  tower.  The  chlorine  is  intro- 
duced from  below  and  dissolves  in  the  liquid.  The  bromine  is  thus 
liberated  and  passes  off  as  vapor.  A  part  of  our  supply  of  bromine  is 
obtained  from  the  brines  of  Ohio,  West  Virginia,  and  Kentucky.  Here 
the  liquid,  after  most  of  the  common  salt  has  been  removed  by  crystal- 
lization, is  assayed  to  ascertain  the  quantity  of  bromine  which  it  con- 
tains, and  is  treated  with  the  calculated  amount  of  sulphuric  acid 
necessary  for  the  action.  Manganese  dioxide  is  then  added  in  small 
quantities  at  a  time.  In  Michigan  the  brines  are  treated  with  elec- 
trolytic chlorine.  The  quantity  produced  in  America  in  1904,  was 
897,100  pounds,  and  was  valued  at  $269,130. 

A  Plan  for  Making  Very  Complex  Equations.  —  When  an  equa- 
tion involves  more  than  two  initial  substances  or  products,  as  does  the 
one  given  above  for  the  first  method  of  preparing  bromine,  it  cannot 
readily  be  worked  out  by  the  method  formerly  recommended  (p.  110). 
After  the  formulae  of  all  the  substances,  on  both  sides,  have  been  set 
down,  it  is  difficult  to  hit  upon  the  proper  numerical  factors  required 
to  balance  the  equation.  In  such  cases  a  good  plan  is  to  select  two  of 
the  initial  substances  and  make  a  partial  equation  showing  part  of  the 
action  and  including  at  least  one  actual  product.  Any  unused  units 
(not  constituting  a  product)  are  then  set  down  also  and  treated  as  a 


THE   HALOGEN   FAMILY  229 

• 
balance.     Thus  the  first  two  of  the  substances,  in  the  above  equation, 

will  furnish  the  second  of  the  products : 

KBr  +  H2S04  -t  KHS04  (+  HBr).  (1) 

Then  we  may  represent  the  formation  of  the  first  of  the  products  from 
the  materials  which  evidently  must  have  been  used  up  in  forming  it, 
and  set  down  the  balance  as  before  : 

Mn02  +  H2S04  -»  MnS04  +  H2O  (+  O).  (2) 

We  then  perceive  that  the  last  of  the  products  might  come  from  the 
oxidation  of  the  first  balance  by  the  second  : 

(2HBr)  -f  (0)  -->  H20  -4-  Br2.  (3) 

The  third  partial  equation  shows  that  2HBr  will  be  needed  for  the 
amount  of  0  obtainable  from  MnOs,  so  we  take  enough  of  the  material 
in  (1)  to  get  the  required  2HBr  : 

2KBr  +  2H2S04  ->  2KHS04  (  +  2HBr).  (1) 

When  we  now  add  the  real  substances  used  and  produced,  as  they 
occur  in  these  partial  equations,  and  leave  out  the  balances,  which 
have  been  adjusted  so  as  to  cancel  one  another,  we  obtain  the  final 
equation  already  given  for  the  action.  It  must  be  observed  that  this 
subdivision  of  the  action  into  parts  is  a  purely  arithmetical  device, 
used  solely  to  simplify  the  arithmetical  process  of  writing  the  equa- 
tions, and  is  not  intended  to  imply  that  the  chemical  change  itself 
follows  these  or,  indeed,  any  stages.  It  happens  that  the  three  par- 
tial equations  we  have  used  in  this  illustration  all  represent  inter- 
actions which  can  take  place  separately.  But  the  arithmetical  value 
of  the  device  does  not  depend  upon  this.  The  partial  equations  made 
for  purposes  like  the  present  one  are  often  purely  fictitious.  It  is 
still  true,  however,  that  we  are  aided  in  the  selection  of  partial  actions 
at  each  step  by  following  some  plausible  theory  as  to  stages  for  the 
action  which,  if  there  were  any,  would  be  chemically  conceivable. 

Physical  Properties.  —  Bromine  is  a  dark-red  liquid  (sp.  gr.  3.18). 
It  boils  at  59°,  forming  a  deep-red  vapor,  and  even  at  ordinary  tem- 
peratures gives  a  high  vapor  pressure  (150  mm.  at  18°)  and  evaporates 
quickly.  When  cooled  it  forms  red,  needle-shaped  crystals,  which  melt 
at -7. 3°.  A  saturated  aqueous  solution  (bromine-water)  at  ordinary 
temperatures  contains  about  three  parts  of  bromine  in  one  hundred  of 


230  INOKGANIC   CHEMISTRY 

water.  The  element  is  much  more  soluble  in  carbon  bisulphide, 
alcohol,  and  other  organic  solvents.  Its  vapor  density  up  to  750°  is  160 
(oxygen  =  32). 

Bromine  (Gk.  j3pufjLo<:,  a  stench)  has  a  most  pungent  odor.  It  has 
a  very  irritating  effect  on  the  mucous  membrane  of  the  nostrils  and 
throat.  If  spilled  upon  the  hands  it  has  a  most  destructive  action 
upon  the  tissues. 

When  free  bromine  is  added  to  starch  emulsion  no  special  change 
in  tint  is  observable  unless  a  very  large  amount  of  the  element  is  used 
(see  Iodine). 

Chemical  Properties.  —  The  molecules  of  bromine  are  much  less 
stable  than  those  of  hydrogen,  oxygen,  or  nitrogen.  The  atomic 
weight  of  bromine  is  79.96,  so  that  in  the  form  of  vapor  it  has  two 
atoms  in  a  molecule,  and  is  represented  by  the  formula  Br2.  At 
1050°  the  vapor  density  is  150.5,  and  dissociation  into  Br  has  begun. 

Like  chlorine,  it  forms  an  unstable  hydrate  with  water. 

Bromine  unites  directly  with  hydrogen.  The  mixture  of  the  gases 
is  not  explosive,  and  the  union  is  much  slower  than  in  the  case  of 
chlorine.  In  presence  of  finely  divided  platinum  the  speed  of  the 
action  may  be  considerably  increased. 

Bromine  forms  compounds  directly,  both  with  non-metals,  like  phos- 
phorus and  arsenic,  and  with  most  of  the  metals.  Towards  unsatu- 
rated  substances  and  organic  compounds  it  behaves  like  chlorine  (q.v.). 
In  all  cases  the  interaction  is  less  violent  than  when  chlorine  is  used, 
and  the  element  is  displaced  from  combination  with  hydrogen  and 
with  the  metals  by  free  chlorine. 

Potassium  bromide  is  employed  in  making  photographic  plates  and 
in  medicine.  Bromine  is  required  in  large  quantities  in  the  manufac- 
ture of  intermediate  products  used  in  the  preparation  of  organic 
dyes. 

HYDROGEN  BROMIDE. 

Preparation.  —  It  might  be  expected  that  the  most  convenient  way 
of  producing  this  compound  would  be  similar  to  that  used  in  prepar- 
ing hydrogen  chloride,  namely,  by  the  action  of  concentrated  sulphuric 
acid  upon  some  common  bromide  such  as  potassium  bromide  (KBr  -f- 
H,S04<=»HBr  +  KHS04).  We  find  indeed  that  at  first  a  colorless 
gas  is  given  off,  which  fumes  strongly  in  the  air  just  like  hydrogen 
chloride,  and  is  the  required  substance.  Almost  immediately,  how- 


THE  HALOGEN  FAMILY 

ever,  this  gas  acquires  a  yellow  and  then  a  brown  tinge,  and  we  dis- 
cover that  free  bromine  is  being  produced  at  the  same  time.  If  we 
examine  the  gas  still  further,  we  recognize  also  the  presence  of  sulphur 
dioxide.  It  is  impossible,  therefore,  to  produce  hydrogen  bromide 
free  from  those  two  impurities  by  this  action. 

The  origin  of  the  bromine  and  sulphur  dioxide  which  complicate 
this  chemical  change  may  readily  be  traced.  Hydrogen  bromide  is 
less  stable  than  hydrogen  chloride,  and  its  hydrogen  can  more  easily 
be  removed  by  the  action  of  substances  containing  oxygen.  In  this 
case  the  sulphuric  acid  acts  as  the  oxidizing  agent,  yielding  oxygen, 
sulphur  dioxide,  and  water  (H2S04  — » 0  +  S02  -f  H20).  Thus  the 
two  extra  gaseous  products  are  seen  to  be  formed  by  a  change  pro- 
ceeding parallel  with  the  main  action : 

2HBr  +  H2S04  -+  2H20  +  SO2  +  Br2. 

The  simultaneous  occurrence,  in  this  fashion,  of  two  more  or  less 
independent  actions  is  not  uncommon.  The  speeds  of  such  actions 
may  be  differently  affected  by  temperature.  Thus,  here,  the  second 
action  seems  to  become  more  extensive  as  the  temperature  rises  (see 
Chap.  xvi).  Since  all  acids  decompose  all  salts  more  or  less,  by  use  of 
an  acid  which  does  not  give  up  its  oxygen  so  readily,  such  as  phos- 
phoric acid,  pure  hydrogen  bromide  may  be  obtained  (KBr  -f-  H3PO4  — > 
HBr  -f-  KH2P04).  The  small  solubility  of  the  salt  in  concentrated 
phosphoric  acid  retards  the  interaction  (p.  179)  and  makes  the  evo- 
lution of  the  gas  very  slow,  however. 

Pure  hydrogen  bromide  is  prepared  by  hydrolysis  (p.  181)  of  phos- 
phorus tribromide.  When  bromine  and  phosphorus  are  mixed,  a  vio- 
lent union  of  the  two  elements  takes  place,  producing  phosphorus 
tribromide  (PBr3).  This  substance,  which  is  a  colorless  liquid,  is  in 
turn  broken  up  with  great  ease  by  water,  producing  phosphorous  acid, 
which  is  not  volatile,  and  hydrogen  bromide  : 

,  Br      HOH 

P  -  Br  +  HOH  ->  P(OH)8  +  3HBr. 
x  Br       HOH 

In  practice,  those  two  actions  are  carried  on  simultaneously.  To  dimin- 
ish the  vigor  of  the  interaction,  red  phosphorus  is  taken  instead  of 
yellow,  and  is  mixed  with  two  or  three  times  its  weight  of  sand  in  a 
flask  (Fig.  70).  A  small  quantity  of  water  is  added.  Excess  of 
water  must  be  avoided,  as  the  hydrogen  bromide  produced  is  extremely 


232 


INORGANIC   CHEMISTRY 


soluble,  and  would  therefore  be  retained  in  the  flask  instead  of  being 
disengaged  as  gas.  The  bromine  is  placed  in  the  dropping  funnel, 

and  admitted,  a  little  at  a 
time,  to  the  mixture.  The 
gas  produced  is  passed 
through  a  U-tube  containing 
glass  beads  mixed  with  red 
phosphorus.  The  latter  com- 
bines with  any  bromine 
which  may  have  escaped 
chemical  change  and  have 
been  carried  along  with  the 
gas.  The  second  U-tube, 
containing  water,  may  be 
attached  when  a  solution  of 
the  gas  is  required. 

FIG  70  Physical     Properties. 

—  Hydrogen    bromide   is   a 

colorless  gas  with  a  sharp  odor.  It  is  two  and  a  half  times  as  heavy  as 
air.  It  is  easily  reduced  to  the  liquid  condition.  It  is  exceedingly  sol- 
uble in  water,  and  in  contact  with  moist  air  condenses  the  water  vapor 
to  clouds  of  liquid  particles.  When  distilled,  the  solution  in  water 
behaves  like  that  of  hydrogen  chloride  (p.  182).  It  loses  mainly  either 
water  or  hydrogen  bromide,  according  as  it  is  dilute  or  exceedingly 
concentrated,  until  an  acid  of  constant  boiling-point  (126°  at  760  mm. 
pressure),  containing  48  per  cent  of  hydrogen  bromide,  passes  over. 

Hydrogen 'bromide,  whether  in  the  gaseous  condition  or  in  the 
liquefied  form,  is  a  nonconductor  of  electricity. 

Chemical  Properties.  —  The  chemical  properties  of  hydrogen 
bromide  are  similar  to  those  of  hydrogen  chloride  (p.  183).  It  is  some- 
what less  stable,  and  dissociation  into  its  constituents  begins  to  be 
noticeable  at  800°.  When  free  from  water,  it  is  not  an  acid  (see 
below),  and  is  not  very  active  chemically,  although  it  behaves  towards 
some  metals  much  like  hydrogen  chloride.  When  the  solution  in 
water  is  strongly  cooled,  crystals  of  a  definite  hydrate  (HBr,  2H20), 
corresponding  to  that  of  hydrogen  chloride,  are  obtained.  When  the 
gas  is  mixed  with  chlorine,  hydrogen  chloride  and  free  bromine  are 
instantly  produced,  and  much  heat  is  evolved  by  the  change,  2HBr-j- 


THE   HALOGEN  FAMILY  233 

C12  — >  2HC1  +  Br2 .  The  heat  produced  by  the  union  of  hydrogen  and 
bromine  vapor  is  12,100  calories.  This  is  much  less  than  the  amount 
produced  by  the  union  of  chemically  equivalent  quantities  of  hydro- 
gen and  chlorine  (22,000  calories).  When  chlorine  displaces  bromine 
from  hydrogen  bromide,  the  heat  evolved  is  found  to  be  the  difference 
between  these  two  numbers.  Using  the  rule  of  constant  heat  summa- 
tion (p.  78),  we  write  equation  (2)  so  that  HBr  is  on  the  same  side 
with  Cl  (with  which  it  interacts),  and  the  products  of  the  equation 
required  (HC1  and  Br)  are  both  on  the  right: 

H  +  Cl     ->  HC1  +  22,000  cal.  (1) 

HBr  -»Br  +  H       -  12,100  cal.  (2) 

Adding,  HBr  +  Cl  -+  HC1  +  Br  +    9900  cal. 

The  12,100  calories  are  produced  by  the  union  of  gaseous  bromine 
with  hydrogen,  and  the  final  result  is,  therefore,  that  for  the  produc- 
tion of  gaseous  bromine.  If  the  heat  of  formation  of  liquid  bromine 
is  required,  the  latent  heat  of  vaporization  of  bromine  (7296  calories), 
which  will  be  evolved  when  the  element  condenses,  must  be  added. 

Chemical  Properties  of  Hydrobromic  Acid. —  The  solution  of 
the  hydrogen  bromide  in  water  is  an  active  acid  (cf.  p.  111).  It  con- 
ducts electricity  extremely  well.  In  contact  with  metals,  oxides  of 
metals,  and  hydroxides  of  metals,  it  behaves  exactly  like  hydrochloric 
acid  (p.  185).  In  the  first  case,  hydrogen  is  set  free  and  the  bromide 
of  the  metal  produced.  In  the  other  two  cases,  water  and  the  bro- 
mides of  the  metals  are  produced.  Oxidizing  agents  set  bromine  free 
from  hydrobromic  acid,  and  the  only  difference  as  compared  with 
hydrochloric  acid  is  that  less  powerful  oxidizing  agents  can  produce 
this  result.  Chlorine  dissolved  in  water  displaces  bromine  from  hydro- 
bromic acid  and  from  soluble  bromides  with  ease. 

IODINE. 

Occurrence. — Iodine,  like  bromine,  occurs  in  sea-water,  although 
it  is  present  in  much  smaller  quantities.  Fortunately,  certain  species 
of  sea- weed  seem  to  have  the  power  to  remove  it  from  water,  and  use 
it  as  a  constituent  in  complex  organic  compounds  which  they  contain. 
In  the  case  of  certain  species,  the  ash  of  the  sea-weed  is  said  to  con- 
tain as  much  as  two  per  cent,  or  even  more.  The  other  chief  source 
of  iodine  is  in  Chili  saltpeter  (NaN03),  in  which  it  is  present  in  the 


234  INORGANIC  CHEMISTRY 

form  of  small  proportions  of  sodium  iodate  (NaI03)  and  sodium 
iodide.  Of  recent  years  the  quantity  obtained  commercially  from  this 
source  has  greatly  increased,  while  that  from  sea-weed  has  diminished. 

It  is  worth  noting  that  while  the  total  amount  of  iodine  found  in  the  human 
body  is  small,  the  proportion  contained  in  the  thyroid  gland  is  consid«rable.  A 
complex  organic  substance,  known  as  iodothyrin,  has  been  extracted  from  sheep's 
thyroids.  It  is  administered  with  marked  success  in  certain  diseases,  such  as 
cretinism,  which  are  associated  with  very  small  development  of  this  gland. 

Preparation.  —  In  factories  where  the  iodine  is  extracted  from 
sea-weed,  the  latter  is  first  burned,  either  roughly  in  hollows  in  the 
ground,  or  more  carefully  in  specially  constructed  ovens.  A  great 
deal  of  the  iodine  seems  to  be  lost  by  volatilization  in  this  process, 
but  the  ash  which  remains,  and  which  is  known  in  Scotland  as  kelp 
and  in  Normandy  as  varec,  still  contains  from  0.5  to  1.5  per  cent  of 
sodium  iodide.  The  ash  is  treated  with  water,  and  the  solution  is  evap- 
orated so  as  to  permit  the  deposition  of  the  sodium  chloride  and  sodium 
sulphate  which  it  contains.  The  sodium  iodide,  being  very  soluble, 
remains  in  the  mother-liquor.  This  is  then  treated  with  manganese 
dioxide  and  sulphuric  acid.  The  quantity  of  manganese  dioxide  is 
carefully  measured  so  as  to  be  just  sufficient  to  set  free  the  iodine 
contained  in  the  liquid  without  proceeding  farther  to  the  liberation  of 
the  chlorine  which  it  contains  in  much  larger  amounts.  When  the 
mixture  is  heated,  the  iodine  passes  off  in  the  form  of  vapor,  and  is 
condensed  in  a  suitable  receiver.  The  action  (cf.  pp.  227,  229)  is  : 

2NaI  +  Mn02  +  3H2SO4  -H>  MnS04  +  2NaHS04  +  2H20  +  I2. 

In  France  the  treatment  is  similar,  excepting  that  chlorine  is  used 
to  liberate  the  iodine  in  the  last  stage  (2NaI  +  C^— »21SraCl  -f  I2). 
The  quantity  is  adjusted  so  that  excess  may  not  be  employed.  This 
is  to  avoid  the  formation  of  a  compound  of  chlorine  and  iodine,  which 
easily  arises  by  direct  union  of  the  elements.  The  iodine,  being  insol- 
uble, forms  a  dense  precipitate,  and,  when  the  liquid  is  pressed  out,  it 
remains  behind  in  the  form  of  a  paste.  Electricity  could  also  be  used 
for  the  decomposition  of  this  mother-liquor.  The  iodine  is  set  free  at 
the  positive  electrode,  while  the  metals  pass  to  the  other.  This 
process  does  not,  however,  seem  as  yet  to  have  found  actual  applica- 
tion in  commerce. 

In  all  cases  the  iodine  is  subjected  to  purification  before  being  sold. 
This  is  carried  out  by  distilling  it  with  a  little  powdered  potassium 


THE   HALOGEN   FAMILY  235 

iodide.  It  condenses  in  the  solid  form  directly,  in  glittering  black 
plates  (sublimed  iodine).  The  distillation  of  a  solid  body,  when  a 
condensation  takes  place  directly  to  the  solid  form,  is  spoken  of  as 
sublimation. 


Physical  Properties.  —  Iodine  (G-k.  louSfe,  like  a  violet)  is  a 
solid  substance  (sp.  gr.  5),  exhibiting  large  crystalline  plates  of  rhombic 
form.  It  melts  at  114°,  and  boils  at  184°.  The  vapor  has  at  first 
a  reddish-violet  tint,  and  on  being  more  strongly  heated  becomes  deep 
blue. 

Iodine  is  much  less  soluble  in  water  than  are  the  other  halogens, 
and  the  solution  has  a  scarcely  perceptible  brown  tint.  At  ordinary 
temperatures  one  part  of  iodine  dissolves  in  about  5000  to  6000  parts 
of  water.  It  is  much  more  soluble  in  carbon  disulphide  (p.  156)  and  in 
chloroform,  in  which  it  gives  violet  solutions.  In  alcohol  it  gives  a 
solution  which  is  brown.  The  brown  color  is  attributed  to  the  fact 
that  in  alcohol  the  iodine  is  in  a  condition  of  feeble  combination  and 
not  simply  in  solution.  An  aqueous  solution  of  potassium  iodide, 
hydrogen  iodide,  or  any  other  iodide,  has  likewise  the  power  to  take  up 
large  quantities  of  iodine.  In  these  cases,  however,  the  formation  of 
definite  compounds  (such  as,  KI  -j-  I2  +±  KI3),  by  a  reversible  action, 
accounts  for  the  amount  of  iodine  which  appears  to  be  dissolved. 

The  behavior  of  free  iodine  towards  starch  forms  a  distinctive  test 
for  both  substances  (cf.  p.  99).  When  starch  is  treated  with  boiling 
water  it  passes  into  a  state  of  fine  suspension,  and  the  liquid  may  be 
filtered  without  removal  of  all  the  starch.  When  traces  of  iodine,  con- 
tained, for  example,  in  the  pale-brown  aqueous  solution,  are  added  to 
this  filtered  starch  emulsion,  a  deep-blue  color  is  produced.  The  same 
action  is  used  as  a  test  for  starch.  This  blue  substance  is  not  a  chemi- 
cal compound,  but  a  solution  of  the  iodine  in  the  solid  starch  which  is 
suspended  in  the  water.  That  the  color  becomes  so  much  stronger  is  in 
harmony  with  the  fact  that  dissolved  bodies  frequently  confer  a  tint 
different  from  their  own  upon  their  solutions. 

Chemical  Properties.  —  The  molecular  weight  of  iodine,  ascer- 
tained by  weighing  the  vapor  at  a  temperature  above  the  boiling-point, 
is  254.8.  The  atomic  weight  being  126.97,  the  molecule  contains  two 
atoms.  This  value  for  the  density  remains  unchanged  when  the 
measurement  is  made  at  temperatures  up  to  about  700°.  Beyond  this 
point,  however,  the  vapor  diminishes  in  density  more  rapidly  than 


236  INORGANIC   CHEMISTRY 

Charles'  law  would  lead  us  to  expect,  and  at  1700°  the  molecular  weight 
has  fallen  to  127.  As  the  vapor  is  heated,  a  larger  and  larger  propor- 
tion of  the  molecules  is  broken  up,  until  the  decomposition  has  become 
complete.  As  in  all  cases  of  dissociation,  when  the  vapor  is  cooled  the 
atoms  recornbine  to  form  molecules.  This  chemical  action  is  interest- 
ing, since  it  is  the  most  notable  case  in  which  we  encounter  both  the 
monatomic  and  the  diatomic  forms  of  the  same  element.  The  heat 
given  out  when  the  atoms  reunite  to  form  the  molecules  is  very  con- 
siderable (21  <=>  I2  +  28,500  cal.),  indicating  that  the  chemical  union 
of  two  atoms  of  identical  nature  may  be  as  vigorous  as  that  of  two 
atoms  of  different  chemical  substances.  The  monatomic  and  diatomic 
forms  of  iodine  should  be  distinct  chemical  substances,  and  if  the 
investigation  of  the  behavior  of  the  former  were  not  hampered  by  the 
very  high  temperature  at  which  alone  it  exists,  it  would  doubtless  be 
found  to  exhibit  different  chemical  properties. 

Iodine  forms  no  hydrate  with  water.  It  exhibits  little  tendency  to 
unite  with  hydrogen.  Iodine  unites  directly  with  a  number  of  ele- 
ments, including  some  non-metals  and  the  majority  of  the  metals. 
When  phosphorus  is  presented  in  the  yellow  form,  the  action  takes 
place  spontaneously  without  the  assistance  of  heat.  Both  chlorine 
and  bromine  displace  iodine  from  combination  with  hydrogen  and  the 
metals  (2HI  +  Br2  —>  2HBr  -f  I2).  The  action  may  be  brought  about 
either  with  the  substances  in  dry  form  or  with  their  aqueous  solutions. 

Iodine  in  water,  like  chlorine  in  water,  constitutes  an  oxidizing 
agent,  although  the  former  is  much  the  feebler  of  the  two.  It  con- 
verts, however,  sulphurous  acid  (q.v.}  into  sulphuric  acid.  The 
action,  so  far  as  the  iodine  and  the  water  are  concerned,  is  shown  by 
the  partial  equation 

I2  +  H,0-»2HI(+0). 

This  action  does  not  take  place,  however,  excepting  in  presence  of  a 
body  capable  of  assisting  in  the  reduction  of  the  water  and  of  uniting 
with  the  oxygen.  In  analytical  chemistry  a  solution  of  iodine  in 
potassium  iodide  containing  a  known  proportion  of  iodine  (a  standard 
solution)  is  used  for  estimating  the  quantity  of  anoxidizable  substance 
present  in  a  given  specimen.  The  amount  of  oxidizable  substance 
present  is  measured  by  the  quantity  of  the  standard  iodine  solution 
which  can  be  decolorized  and  suffer  removal  of  its  iodine.  This 
method  is  known  as  iodimetry. 

Iodine  and  its  compounds  are  much  used  in  the  arts  and  medicine. 


THE   HALOGEN  FAMILY  237 

Iodine  is  applied,  in  the  form  of  an  alcoholic  solution  ("tincture  of 
iodine"),  for  the  reduction  of  some  swellings.  It  is  required  in  mak- 
ing iodoform  (CHI3),  and  the  iodides  of  potassium,  rubidium,  and 
sodium,  which  are  used  in  medicine.  Potassium  iodide  is  likewise  em- 
ployed in  the  manufacture  of  photographic  plates.  Large  quantities  of 
iodine  find  application  in  the  color  industry  for  making  intermediate 
products  needed  in  the  manufacture  of  aniline  and  other  organic  dyes. 

HYDROGEN  IODIDE. 

Preparation. — The  direct  union  of  hydrogen  and  iodine  cannot  be 
employed  in  preparing  pure  hydrogen  iodide.  The  union  takes  place 
slowly,  and,  since  the  action  is  markedly  reversible  (cf.  p.  176),  always 
remains  incomplete  :  H2  -f-  I2  +±  2HI.  In  presence  of  finely  divided 
platinum  the  union,  so  far  as  it  is  capable  of  proceeding,  may  be 
hastened,  but  the  amount  of  the  compound  formed  is  not  greater.  At 
448°,  for  example,  79  per  cent  of  the  constituents  unite,  and  the 
remainder  is  unaffected  either  by  longer  heating  or  by  platinum.  At 
other  temperatures  the  proportion  is  different,  but  at  none  is  the 
action  complete.  Hydrogen  and  iodine  unite  with  little  vigor  because 
the  union  is  not  accompanied  by  a  sufficiently  marked  liberation  of 
energy  (cf.  HC1,  p.  233).  Thus  at  18°  the  action  is  really  endothermal 
(—  6100  calories  for  HI),  and  at  400°,  although  the  action  is  exother- 
mal, only  535  calories  of  heat  are  set  free. 

The  action  of  concentrated  sulphuric  acid  upon  potassium  or  sodium 
iodide  is  equally  inapplicable.  In  this  case,  as  in  that  of  hydrogen 
bromide  (p.  231),  the  hydrogen  halide  reduces  the  sulphuric  acid,  and 
much  free  iodine  is  formed.  Here,  on  account  of  the  greater  instabil- 
ity of  hydrogen  iodide  and  the  greater  ease  with  which  it  parts  with 
its  hydrogen,  the  reduction  of  the  sulphuric  acid  is  much  more  com- 
plete, the  primary  product  being  apparently  hydrogen  sulphide.  The 
actions,  which  proceed  simultaneously  (p.  231),  are: 

KI  +  H2S04^±  HI  +  KHS04  andH2S04  +  8HI 1*  H2S  +  4H20  +  4Ir 

As  soon  as  the  heat  produced  by  the  action  has  raised  the  temperature 
sufficiently,  hardly  any  of  the  hydrogen  iodide  escapes  oxidation.* 

*  When  much  sulphuric  acid  is  used,  sulphur  dioxide  and  free  sulphur  are 
formed  also.  This  is  in  consequence  of  a  secondary  action  of  the  hydrogen  sulphide 
on  the  sulphuric  acid,  and  of  the  sulphur  dioxide  so  formed  upon  the  excess  of 
hydrogen  sulphide  (see  Hydrogen  sulphide)  : 


238  INORGANIC  CHEMISTRY 

Finely  powdered  sodium  iodide  and  concentrated  phosphoric  acid 
(cf.  p.  179),  when  mixed  and  warmed,  give  pure  hydrogen  iodide 
(Nal  +  H3P04<i±HI  J  -f-  NaltjPO*).  This  action  was  formerly  used 
in  preparing  the  gas. 

The  best  method  is  one  similar  to  that  described  under  hydrogen 
bromide.  Phosphorus  and  iodine  unite  directly  to  form  PI8.  This  is 
a  yellow  solid  which  is  violently  decomposed  by  water  and  gives  phos- 
phorous acid  and  hydrogen  iodide  : 

PI8  +  3H/)  ->  P(OH),  +  SHI. 

If  excess  of  water,  which  dissolves  hydrogen  iodide,  is  avoided,  the 
latter  goes  off  in  a  continuous  stream  in  a  gaseous  condition. 

Still  another  method  of  making  hydrogen  iodide  is  frequently 
employed  when  a  solution  of  the  gas  in  water  is  required,  and  not  the 
gas  itself.  Powdered  iodine  is  suspended  in  water,  and  hydrogen  sul- 
phide gas  (q.v.)  is  introduced  through  a  tube  in  a  continuous  stream. 
The  iodine  dissolves  slowly  in  the  water,  I2  (solid)  <=±  I2  (diss'd),  and 
acts  upon  the  hydrogen  sulphide,  which  likewise  dissolves,  H2S  (gas)  +± 
H2S  (diss'd).  Sulphur  separates  in  a  fine  powder,  S  (diss'd)  <=±  S 
(solid),  and  hydrogen  iodide  is  formed  in  accordance  with  the 
equation  : 

S. 


This  action  takes  place,  however,  only  in  presence  of  water,  although 
the  water  does  not  appear  in  the  equation.  The  solution  is  freed  from 
the  deposit  of  sulphur  by  filtration,  and  may  be  concentrated  to  57 
per  cent  of  hydriodic  acid  by  distilling  off  the  water. 

The  theory  of  the  last  method  is  worthy  of  attention.  It  will  be 
seen  that  while  iodine  has  little  tendency  to  unite  with  free  hydrogen, 
it  is  here  able  to  decompose  a  compound  containing  hydrogen,  in  order 
to  secure  this  element.  It  is  enabled  to  do  this  by  the  fact  that  the 
very  large  amount  of  heat  given  out  by  the  mere  solution  of  hydrogen 
iodide  in  water  converts  the  action,  which  would  otherwise  be  endother- 
mal,  into  an  exothermal  one.  In  the  absence  of  water,  the  reverse  of 
the  above  action  takes  place  with  ease.  In  presence  of  water,  however, 
the  great  heat  of  solution  (p.  164)  of  the  hydrogen  iodide  (  +  19,- 
200  cal.)  more  than  balances  the  heat  absorbed  by  the  chemical  change, 
and  the  action  as  a  whole  takes  place  with  evolution  of  heat  (see,  also, 
Preparation  of  hydrogen  sulphide). 


THE   HALOGEN  FAMILY  239 

Physical  Properties.  —  Hydrogen  iodide  is  a  colorless  gas  with 
a  sharp  odor.  Its  molecular  weight  is  128,  and  it  is  therefore  much 
heavier  than  air,  the  average  weight  of  whose  molecules  is  28.955 
(p.  193).  It  is  a  nonconductor  of  electricity,  both  in  the  gaseous 
and  in  the  liquefied  conditions.  It  is  exceedingly  soluble  in  water,  so 
that  at  0°  ten  grams  of  water  will  absorb  ninety  grams  of  the  gas,  giv- 
ing a  90  per  cent  solution.  The  behavior  of  this  solution  is  similar  to 
those  of  hydrogen  chloride  and  hydrogen  bromide  (cf.  pp.  182,  232). 
The  mixture  of  constant  boiling-point  distils  over  at  127°  (at  760  mm.), 
and  contains  57  per  cent  of  hydrogen  iodide. 

Chemical  Properties*  —  Hydrogen  iodide  Is  the  least  stable  of  the 
hydrogen  halides.  When  heated  it  begins  visibly  to  decompose  into 
its  constituents  at  180°.  On  account  of  the  ease  with  which  it  parts 
with  the  hydrogen  which  it  contains,  it  can  be  burned  in  oxygen  gas, 
4HI  -J-  O2  — >  2H20  +  2I2.  When  the  gas  is  mixed  with  chlorine,  a 
violent  chemical  change,  accompanied  by  a  flash  of  light,  occurs,  the 
iodine  is  set  free,  and  hydrogen  chloride  is  produced,  CLj  +  2HI— -> 
2HC1  -f  I2.  Bromine  vapor  will  similarly  displace  the  iodine  from 
hydrogen  iodide. 

Chemical  Properties  of  Hydriodic  Acid.  —  In  most  respects 
the  aqueous  solution  behaves  exactly  like  hydrochloric  and  hydro- 
bromic  acid.  With  oxidizing  agents,  for  example,  such  as  manganese 
dioxide,  it  gives  free  iodine,  just  as  the  others  give  free  chlorine  and 
bromine  respectively.  Here,  however,  the  oxidation  is  so  much  more 
easily  carried  out,  that  it  is  slowly  effected  by  atmospheric  oxygen, 
so  that  hydriodic  acid  left  exposed  to  the  air  gradually  becomes  brown 
(02  4-  4HI  ->  2  H2O  +  2I2).  The  free  iodine  which  is  produced,  re- 
mains dissolved  in  the  hydrogen  iodide,  probably  in  the  form  of  a 
compound  HI8.  Finally,  however,  the  free  iodine  as  its  quantity 
becomes  greater,  and  that  of  the  hydrogen  iodide  smaller,  is  deposited 
in  crystalline  condition.  On  account  of  the  ease  with  which  hydriodic 
acid  parts  with  its  hydrogen,  it  is  frequently  used  in  chemistry  as  a 
reducing  agent. 

Although  the  dry  gas  is  not  an  acid,  the  solution  has  all  the  ordi- 
nary properties  of  this  class  of  substances  (cf.  p.  92).  The  hydrogen 
may  be  displaced  by  metals  like  zinc  and  magnesium.  The  acid  inter- 
acts with  oxides  and  hydroxides,  forming  iodides  and  water. 


240 


INORGANIC   CHEMISTRY 


FLUORINE. 

The  discussion  of  this  element  should  logically  have  preceded  that 
of  chlorine,  since  it  is  of  all  the  members  of  the  halogen  family  the 
most  active.  Chlorine  was  taken  up  first,  however,  because  its  com- 
pounds are  more  familiar.  Fluorine  is  found  in  combination  in  nature. 
It  occurs  chiefly  in  the  mineral  fluorite  (calcium  fluoride,  CaF2)  and  in 
cryolite,  a  double  fluoride  of  aluminium  and  sodium  (A1F3,  3NaF). 


Preparation.  —  When  a  solution  of  hydrofluoric  acid  is  heated 
with  manganese  dioxide,  oxidation  does  not  occur  and  free  fluorine  is 
not  produced.  Until  very  recently  all  efforts  to  isolate  the  element 
failed.  It  was  perfectly  understood  that  the  reason  of  these  failures 

lay  in  the  greater  chemical  activity  of  fluor- 
ine, which  made  it  more  difficult  of  separa- 
tion from  any  state  of  combination  than 
the  other  halogens.  Its  preparation  was 
finally  achieved  by  Moissan  (1886)  by  the 
decomposition  of  anhydrous  hydrogen  fluor- 
ide, which  is  liquid  below  19°,  by  means 
of  electricity.  The  apparatus  (Fig.  71)  is 
made  of  copper,  which,  after  receiving  a 
thin  coating  of  the  fluoride,  is  not  further 
affected.  To  reduce  the  tendency  to  chem- 
ical union,  the  whole  is  immersed  in  a  bath 
giving  a  temperature  of  —  23°.  The  elec- 
trodes are  made  of  an  alloy  of  platinum  and 
iridium,  which  is  the  only  substance  that 
can  resist  the  action  of  the  fluorine  when 
freshly  liberated  by  the  electric  current. 
Hydrogen  fluoride,  like  other  hydrogen 
halides,  is  a  nonconductor  of  electricity,  and 

a  small  quantity  of  potassium  fluoride  has  to  be  added  to  enable  the 
current  of  electricity  to  pass.  The  fluorine  is  set  free  at  the  positive 
electrode,  and  hydrogen  appears  at  the  negative.  The  U-tube  is  closed 
after  the  introduction  of  the  hydrogen  fluoride  by  means  of  blocks  made 
of  calcium  fluoride,  which  is  naturally  unable  further  to  enter  into  com- 
bination with  fluorine.  For  the  reception  and  examination  of  the 
fluorine  gas,  other  copper  tubes  can  be  screwed  on  to  the  side  necks  of 
the  apparatus,  and,  when  necessary,  small  windows  of  calcium  fluoride 


FIG.  71. 


THE  HALOGEN  FAMILY  241 

can  be  provided.  It  has  been  found  that  fluorine  dried  with  extraor- 
dinary precautions  is  without  action  on  glass. 

Physical  Properties.  —  Fluorine  is  a  gas  whose  color  is  like  that 
of  chlorine,  but  somewhat  paler.  Its  density  has  not  been  measured 
with  great  exactitude,  but  the  value  obtained  indicates  a  molecular 
weight  of  38,  showing  that  there  are  two  atoms  in  the  molecule  (the 
atomic  weight  is  19).  The  gas  is  the  most  difficult  of  the  halogens  to 
liquefy.  The  liquid  boils  at  -  186°. 

Chemical  Properties.  —  Fluorine  unites  with  every  element,  with 
the  exception  of  oxygen,  and  in  many  cases  does  so  with  such  vigor 
that  the  union  begins  spontaneously  without  the  assistance  of  external 
heat.  Dry  platinum  and  gold  are  the  elements  least  affected.  It 
explodes  with  hydrogen  at  the  ordinary  temperature,  without  the 
assistance  of  sunlight.  Fluorine  displaces  oxygen  from  water  instan- 
taneously and  gives  ozone  (q-v.).  On  the  introduction  of  a  drop  of 
water  into  a  tube  of  fluorine  the  vessel  is  filled  with  the  deep-blue 
gas,  3F2  +  3H2O  -»  3H2F2  +  08. 

The  chlorine  in  hydrogen  chloride  is  displaced  by  fluorine  as  easily 
as  chlorine  in  turn  displaces  bromine  or  iodine. 

HYDROGEN  FLUORIDE. 

Preparation.  —  Pure,  dry  hydrogen  fluoride  is  best  made  by  heat- 
ing potassium  hydrogen,  fluoride,  2KHF2  +±  2KF  -f  H2F2  f .  For 
ordinary  purposes,  however,  the  preparation  of  an  aqueous  solution  is 
the  ultimate  object.  Usually  powdered  calcium  fluoride  is  treated 
with  concentrated  sulphuric  acid,  and  the  mixture  distilled  in  a  plat- 
inum retort : 

CaF2  +  H2S04  <p±  CaS04  -f  HaF2 1 . 

The  hydrofluoric  acid  passes  over  and  is  caught  in  distilled  water. 
The  aqueous  solution  thus  obtained  has  to  be  kept  in  vessels  made  of 
lead,  rubber,  or  paraffin,  as  glass  interacts  with  the  acid  with  great 
rapidity  (see  below). 

Physical  Properties.  —  Hydrogen  fluoride  is  a  colorless  liquid, 
boiling  at  19.4°.  It  mixes  freely  with  water,  and,  on  distillation,  an 
acid  of  constant  boiling-point  (120°  at  760  mm.)  containing  35  per 
cent  of  hydrogen  fluoride  is  obtained.  The  vapor  density  of  the 


242  INORGANIC   CHEMISTRY 

hydrogen  fluoride  between  its  boiling-point  and  30°  corresponds  to  a 
molecular  weight  of  40,  and  the  formula  should  therefore  be  H2F2. 
Above  this  temperature  the  vapor  becomes  lighter  (p.  207),  and,  when 
88°  is  reached,  the  molecular  weight  has  fallen  to  20,  corresponding 
to  the  formula  HF. 

Association.  —  Many  compounds  resemble  hydrogen  fluoride  in 
seeming  to  consist  of  molecules  which  are  multiples  of  the  simplest 
possible.  This  is  spoken  of  as  indicating  a  tendency  to  association.* 
Thus,  sulphuric  acid  and  nitric  acid  in  the  liquid  condition  are  com- 
posed of  more  complex  aggregates  than  H2S04  and  HN08.  Even  water 
is  largely  (H20)2  or  even  (H20)8,  although  the  vapor  is  H2O.  In  such 
cases  dissociation  into  the  simpler  molecules  takes  place  gradually  as 
the  temperature  is  raised.  Sometimes  this  doubling  up  of  molecules 
is  called  polymerization. 

In  cases  of  association  the  observed  molar  weights,  at  low  tem- 
perature, are  multiples  of  the  smallest  possible  molar  weight  required 
for  the  simplest  formula  (in  above  instance,  HF).  But  many  sub- 
stances naturally  possess  formulae  which  are  multiples  of  the  simplest 
without  showing,  as  the  temperature  is  raised,  any  tendency  to  pro- 
gressive dissociation  into  the  corresponding  simplest  molecules.  Thus, 
acetylene  (p.  204)  is  C2H2  at  all  temperatures,  and  acetic  acid  (g.v.), 
although  it  is  associated  to  C4H804  at  its  boiling-point,  never  becomes 
simpler  than  C2H4O2  at  any  temperature. 

Chemical  Properties  of  Hydrofluoric  Acid.  —  Metals  like  zinc 
and  magnesium  interact  with  hydrofluoric  acid  with  evolution  of 
hydrogen.  The  action  is  less  violent  than  with  other  halogen  acids. 
The  acid  interacts  with  oxides  and  hydroxides,  forming  fluorides. 
The  chief  difference  in  this  respect  which  it  exhibits,  when  compared 
with  the  other  halogen  acids,  is  one  which  we  should  expect  from  its 
formula,  H2F2.  We  may  displace  either  one  or  both  the  hydrogen 
atoms  in  the  molecule  with  a  metal.  Thus,  one  of  the  commonest  salts 
of  hydrofluoric  acid  is  potassium  hydrogen  fluoride  KHF2  mentioned 
above,  KOH  +  H2F2  -» KHF2  +  H20.  In  this  respect  the  acid  re- 
sembles sulphuric  acid  and  other  acids  containing  more  than  one 
replaceable  hydrogen  unit.  Salts  in  which  a  portion  of  the  acid 
hydrogen  still  remains  undisplaced  are  spoken  of  as  acid  salts. 

*  The  elementary  substances  show  similar  behavior  (pp.  205,  206). 


THE   HALOGEN   FAMILY  243 

The  most  remarkable  property  of  hydrofluoric  acid  depends  on 
the  great  tendency  which  fluorine  has  to  unite  with  silicon,  forming 
the  gaseous  silicon  tetrafluoride.  Glass,  which  is  commonly  made 
by  fusing  together  sodium  carbonate,  calcium  oxide,  and  sand  (silicon 
dioxide),  is  a  mixture  of  silicates  of  calcium  and  sodium,"  and  is 
rapidly  decomposed  by  hydrofluoric  acid.  The  nature  of  the  change 
is  shown  by  the  two  following  equations  : 

CaSi08   +3H2F2-+  SiF4  -f-  CaF2   +  3H20, 
Na2Si08  +3H2F2-»  SiF4  +  2NaF  +  3H2O. 

All  other  silicates  are  decomposed  according  to  the  same  plan.  The 
silicon  tetrafluoride  passes  off.  The  fluorides  of  calcium  and  sodium 
are  solid  and  crumble  away  or  dissolve.  Thus  the  glass  is  completely 
disintegrated.  The  vapor  of  hydrofluoric  acid,  generated  in  the  way 
described  above  from  calcium  fluoride  in  a  lead  dish,  is  used  for  etch- 
ing glass.  The  surface  of  the  glass  is  covered  with  paraffin  to  protect 
it  from  the  action  of  the  vapor,  and  with  a  sharp  instrument  portions 
of  this  paraffin  are  removed  where  the  etching  effect  is  desired.  The 
vapor  gives  a  rough  surface  where  it  encounters  the  glass.  The 
aqueous  solution,  which  may  also  be  employed,  makes  smooth  depres- 
sions on  the  surface.  The  acid  is  also  used  in  the  analysis  of  minerals 
containing  silicates,  which  frequently  are  not  attacked  by  other  acids. 

THE  HALOGENS  AS  A  FAMILY. 

It  may  be  useful  here  to  bring  together  some  of  the  facts  in  regard 
to  the  halogens  and  their  compounds  by  way  of  showing  more  clearly 
how  far  they  resemble  one  another,  and  in  what  ways  they  differ. 
The  most  noticeable  fact  is  that,  if  we  arrange  them  in  order  in  respect 
to  any  one  property  chemical  or  physical,  the  other  properties  -will 
be  found  to  place  them  in  the  same  order.  Thus,  if  we  consider 
(1)  the  physical  properties,  we  find  that  the  color  deepens  as  we  pass 
from  fluorine  through  chlorine  and  bromine  to  iodine.  The  specific 
gravities  of  the  elements  increase  in  the  same  order.  The  volatility 
of  the  elements  decreases  in  the  same  way  —  fluorine  being  the  hard- 
est to  liquefy,  while  iodine  is  a  solid  and  boils  at  a  fairly  high  tem- 
perature. (2)  In  their  chemical  behavior,  when,  for  example,  they 
unite  with  the  metals  and  hydrogen,  the  vigor  of  the  action  is  greatest 
with  fluorine  and  diminishes  progressively  until  we  reach  iodine. 
We  shall  see  later  that  the  affinity  for  oxygen,  on  the  other  hand, 
increases  as  we  pass  from  fluorine  to  iodine. 


244  INORGANIC    CHEMISTRY 

(3)  The  relations  of  these  elements  in  combination  show  that  they 
are  all  univalent  in  respect  to  union  with  hydrogen  and  metals.  In 
their  oxygen  compounds  (q.v.),  however,  they  frequently  exhibit  a 
higher  valence.  The  compounds  which  they  form  with  any  one  ele- 
ment are  usually  very  similar  to  one  another.  All  the  hydrogen  com- 
pounds, for  example,  become  acids  when  dissolved  in  water.  The  most 
noticeable  lack  of  harmony  in  this  group  is  observed  when  we-  consider 
the  solubilities  of  the  corresponding  compounds.  Thus,  the  potassium 
salts  are  all  soluble  in  water.  Silver  chloride,  bromide,  and  iodide 
are  almost  insoluble,  the  amount  dissolved  decreasing  in  that  order. 
Silver  fluoride,  however,  is  quite  soluble.  Calcium  chloride,  bromide, 
and  iodide  are  all  very  soluble,  while  calcium  fluoride  is  almost  com- 
pletely insoluble. 

It  will  be  noted  that  the  order  in  which  the  elements  are  thus 
placed  by  consideration  of  most  of  their  properties  is  the  order  of 
increasing  atomic  weights  (see  Periodic  system). 

COMPOUNDS  OF  THE  HALOGENS  WITH  EACH  OTHER. 

We  have  incidentally  mentioned  the  fact  that  iodine  unites  with 
chlorine  to  form  a  definite  compound.  In  reality  there  are  two  such 
compounds.  The  most  familiar  is  a  red  crystalline  substance  having 
the  composition  IC1.  Another  compound,  IC18,  is  made  by  the  use 
of  excess  of  chlorine.  Iodine  unites  with  bromine  to  form  the  com- 
pound IBr,  while  a  compound  with  fluorine,  to  which  the  composition 
IF6  has  been  assigned,  is  supposed  to  exist.  None  of  these  compounds 
are  particularly  stable,  and  some  of  them  decompose  very  easily.  It 
is  frequently  stated  that  elements  which  resemble  one  another  chemi- 
cally show  little  tendency  to  chemical  union.  Yet  in  the  case  of 
IBr,  for  example,  the  tendency  to  decompose  into  the  elements 
(2IBr  — » I2  -f-  Br2)  must  be  interpreted  as  meaning  that  the  iodine 
and  bromine  prefer  to  unite  with  themselves  to  form  the  molecules 
I2  and  Br2  rather  than  with  one  another.  In  view  of  this  the  above 
remark  loses  some  of  its  point,  for  an  element  certainly  resembles  itself 
more  than  it  does  any  other,  and  the  compounds  CLj,  H2,  etc.,  are 
amongst  the  most  stable  that  we  know. 

Exercises.  —  1.  What  impurities  is  commercial  iodine  likely  to 
contain  ?  In  what  way  does  heating  with  potassium  iodide  (p.  234) 
free  it  from  these  ? 


THE   HALOGEN  FAMILY  245 

2.  Classify  all  the  chemical  actions  in  this  chapter  according  as 
they  belong  to  one  or  other  of  the  ten  kinds  (p.  187). 

3.  Explain  the  condensation  of  atmospheric  moisture  by  the  hydro- 
gen halides  (cf.  p.  162). 

4.  Tabulate,  more  fully  and  specifically  than  is  done  in  the  section 
on  "The  Halogens  as  a  Family,"  (a)  the  physical  properties,  (£)  the 
chemical  properties,  (c)  the   chemical   relations,  of   the  members  of 
this  group. 

5.  Construct  the  equation  on  p.  234  by  the  use  of  partial  equations 
as  in  the  example  on  p.  229. 

6.  Using  the  method  given  on  p.  229,  construct  a  single  equation 
for  the  formation  of  iodine,  water,  and  hydrogen  sulphide  directly  from 
potassium  iodide  and  sulphuric  acid. 


CHAPTER   XV 
CHEMICAL   EQUILIBRIUM 

IN  spite  of  its  formidable  title,  this  chapter  will  introduce  nothing 
novel.  Its  purpose  is  to  collect  together  and  organize  more  definitely 
a  number  of  scattered  facts  and  ideas  which  have  already  come  up  in 
various  connections.  On  this  account,  however,  it  will  be  all  the  more 
necessary  for  the  reader  to  refresh  his  remembrance  of  these  facts 
and  ideas  by  re-reading  the  pages  to  which  reference  may  be  made. 

Reversible  Actions.  —  In  discussing  the  union  of  hydrogen  and 
iodine  (p.  237),  it  was  stated  that  the  progress  of  the  action  ceases 
while  yet  a  large  amount  of  both  the  substances  necessary  for  its  main- 
tenance still  remains  available.  Now  the  materials  left  over  are  pre- 
sumably no  less  capable  of  uniting  than  the  parts  which  have  already 
united.  The  solution  of  this  mystery  lies  in  the  fact  (p.  239)  that 
decomposition  of  the  compound  can.  begin  at  180°,  and  therefore  takes 
place  actively  at  448°.  Hence  the  product  of  the  union  must  begin 
to  dissociate,  in  part  at  least,  as  soon  as  any  of  it  is  formed.  Thus 
two  changes,  one  of  which  undoes  the  work  of  the  other,  must  go  on 
simultaneously.  In  consequence  of  this,  neither  can  reach  comple- 
tion. As  we  should  expect,  experiment  shows  that  it  makes  no  differ- 
ence whether  we  start  with  the  elements  or  with  the  compound  :  the 
proportions  of  the  materials  found  in  the  tube,  after  it  has  been  heated 
for  a  sufficient  length  of  time,  are  in  both  cases  the  same.  A  general 
statement  may  be  founded  on  facts  like  this,  to  the  effect  that  a  chem- 
ical action  must  remain  more  or  less  incomplete  when  the  reverse 
action  also  takes  place  under  the  same  conditions  (cf.  p.  35).  Two 
arrows  pointing  in  opposite  directions  are  used  in  equations  repre- 
senting reversible  changes :  * 

H2  +  I2<=±  2HI  or  2HI  <=>  H2  +  I,. 

It  will  be  observed  that  representing  reversible  actions  by  equations  involves 
a  departure  from  the  original  meaning  of  an  equation.  Thus  at  448°,  80  per  cent 

*  The  reader  must  avoid  the  idea  that  a  reversible  action  is  one  which  goes  to 
completion,  and  then  runs  back  to  a  certain  extent.  This  conception  would  be 
contrary  to  the  fact,  and  opposed  to  the  principles  of  energetics,  as  well  as  inex- 
plicable by  the  kinetic  hypothesis. 


CHEMICAL  EQUILIBRIUM  247 

of  the  substances  are  in  the  form  HI  and  20  per  cent  in  the  uncombined  state: 


In  other  words,  the  amounts  of  matter  on  the  two  sides  are  not  equal.  Each  side, 
taken  separately,  shows  correctly  the  proportions  used  in  the  interaction  for  which 
it  stands,  however.  Hence  the  equation  in  a  reversible  action  professes  to  show 
quantitatively  the  change  which  would  occur  if  each  of  the  two  opposed  actions  it 
includes  were  to  be  allowed  separately  to  proceed  to  completion. 

The  following  are  some  other  examples  of  actions  of  the  same 
kind:  (1)  The  interactions  of  sulphuric  acid  and  sodium  chloride 
(p.  179),  and  (2)  of  chlorine  and  water  (p.  176),  which  were  fully 
discussed  at  the  time;  (3)  that  of  equivalent  amounts  of  iron  and 
water,  or  of  magnetic  oxide  of  iron  and  hydrogen  (p.  110),  which  does 
not  proceed  to  completion  in  either  direction  when  either  set  of  mate- 
rials is  sealed  up  in  a  tube  and  heated  ;  (4)  the  behavior  of  barium 
peroxide  (p.  63),  of  hydrates  (p.  120),  of  iodine  vapor  (p.  205),  of 
water  vapor  (p.  119),  of  sulphur  vapor  (p.  205),  of  phosphorus  vapor 
(p.  205),  and  of  mercuric  oxide  (pp.  62,  77). 

When  the  action  is  one  which  is  reversible,  but,  under  the  circum- 
stances being  discussed,  proceeds  far  towards  completion  in  one  direc- 
tion, the  arrow  will  be  modified  to  indicate  this  fact: 
C12  +  H20  <=;  HC1  +  HC10  (p.  176). 

When  this  relative  completeness  is  due  to  precipitation  or  volatiliza- 
tion, the  fact  will  be  indicated  by  vertical  arrows  : 

KaCl  +  H2S04fc?NaHS04  +  HC1  f  (p.  178). 
NaClJ+  H2S04±+NaHS04  +  HC1  (p.  179). 

All  chemical  actions  do  not  belong  to  this  class.  Many  proceed 
uninterruptedly  to  exhaustion  of  one,  or  all,  of  the  ingredients.  For 
example,  equivalent  amounts  of  magnesium  and  oxygen  combine  com- 
pletely (2Mg  4-  02  —  >  2MgO).  Here,  however,  the  product  is  not 
decomposed  even  at  the  white  heat  produced  by  the  vigor  of  the  union. 
Indeed,  magnesium  oxide  cannot  be  decomposed,  and  the  action  re- 
versed, at  any  temperature  we  can  command.  The  other  complete 
actions  are  so  because  they  are  likewise  irreversible. 

Chemists  now  look  upon  reversibility  as  being  the  normal  property  of  all  chem- 
ical changes.  They  consider  the  lack  of  obvious  incompleteness  as  being  due  to 
the  very  small  degree  to  which  the  reverse  action  can  occur  under  the  given  con- 
ditions, rather  than  to  the  entire  absence  of  any  reverse  action. 

Kinetic  Explanation.  —  Restating  these  facts  in  terms  of  the 
Jkinetic  hypothesis  will  enable  us  to  reason  more  clearly  about  this 


248  INORGANIC  CHEMISTRY 

variety  of  chemical  change.  Suppose  we  start  with  the  materials  rep- 
resented on  one  side  of  such  an  equation.  The  molecules  of  these 
materials  will  encounter  one  another  frequently  in  the  course  of  their 
movements.  In  a  certain  proportion  of  these  collisions  the  chemical 
change  will  take  place.  In  the  earliest  stages  there  will  be  few  of  the 
new  kind  of  molecules,  but,  as  the  action  goes  on,  these  will  increase 
in  quantity.  There  will  be  two  consequences  of  this.  In  the  first 
place  the  parent  materials  will  diminish  in  amount,  the  collisions 
between  their  molecules  will  become  fewer,  and  the  speed  of  the  for- 
ward action  will  therefore  become  less  and  less.  In  the  second  place 
the  increase  in  the  number  of  molecules  of  the  products  will  result  in 
more  frequent  collisions  between  them,  in  more  frequent  occurrence  of 
the  chemical  change  which  they  can  undergo,  and  thus  in  an  increase 
in  the  speed  of  the  reverse  action.  The  forward  action  decreases  in 
speed  progressively ;  the  reverse  action  increases  in  speed.  Finally 
the  two  speeds  must  become  equal,  and  at  that  point  perceptible 
change  in  the  condition  of  the  whole  must  cease. 

^The  most  immediate  inference  from  this  mode  of  viewing  the 
matter  is,  that  the  apparent  halt  in  the  progress  of  the  action  does 
not  indicate  any  cessation  of  either  chemical  change.  Both  changes 
must  go  on  in  consequence  of  the  continued  encounter  of  the  proper 
molecules.  But  since  they  proceed  with  equal  speed  they  produce  no 
alteration  in  the  mass  as  a  whole.  In  fact,  the  final  state  is  one  of 
equilibrium,  and  not  of  repose.  Hence,  chemical  changes  which  are 
reversible  lead  to  that  condition  of  seemingly  suspended  action  which 
we  speak  of  as  chemical  equilibrium.  The  changes  themselves  are 
called  reversible,  or,  since  they  represent  a  state  of  balance  between 
opposing  tendencies,  balanced  actions. 

The  detailed  discussion  of  the  relations  of  liquid  and  vapor  (pp.  117, 
135),  of  ice  and  water  (p.  115),  and  of  saturated  solution  and  undis- 
solved  solid  (pp.  152,  160),  liquid  or  gas  (p.  154),  has  already  famil- 
iarized us  with  this  term  and  its  significance.  By  the  use  of  the 
kinetic  hypothesis  we  can,  in  fact,  apply  the  very  same  sets  of  ideas 
to  the  discussion  of  any  kind  of  reversible  phenomena. 

When  this  is  done  for  reversible  chemical  actions  we  perceive  that 
two  circumstances,  in  particular,  are  likely  to  have  a  noteworthy  effect 
in  the  promotion  or  retardation  of  a  given  chemical  change.  These 
are  the  homogeneity,  or  the  reverse,  of  the  mixture,  and  the  molecular 
concentration  of  each  component. 


CHEMICAL   EQUILIBRIUM  249 

The  Influence  of  Homogeneous  Mixture.  —  Evidently  the  inter- 
action of  substances  will  be  greatly  favored  if  they  are  capable  of 
remaining  intimately  mixed.  When  this  is  the  case,  as  in  gaseous 
or  liquid  mixtures,  every  molecule  has  an  equal  opportunity  freely  to 
encounter  every  other  molecule.  On  the  other  hand,  when  one  sub- 
stance is  a  liquid  and  the  other  a  gas  not  especially  soluble  in  the 
liquid,  or  when  one  is  a  liquid  and  the  other  a  solid  suspended,  but 
not  dissolved,  in  the  liquid,  or  when  they  are  solid  and  gas  respec- 
tively, the  chances  for  mutual  encounters  between  the  two  different 
kinds  of  molecules  will  be  very  notably  restricted.  Thus  the  progress 
of  an  action  between  inhomogeueously  mixed  bodies,  otherwise  capable 
of  interaction,  must  be  greatly  affected  by  purely  physical  circum- 
stances. Some  of  the  remarkable  consequences  of  this  were  discussed 
fully  in  connection  with  the  preparation  of  hydrogen  chloride  (p.  179). 

The  Influence  of  Molecular  Concentration.  —  Even  when  the 
mixture  is  homogeneous,  a  second  factor  will  affect  the  action.  The 
frequency  of  the  encounters  amongst  a  given  set  of  molecules,  result- 
ing in  a  definite  chemical  change,  will  evidently  depend  entirely  upon 
the  degree  to  which  they  are  concentrated  in  each  other's  neighborhood. 
Larger  amounts  of  one  of  the  materials,  for  example,  will  not  result 
in  more  rapid  chemical  action  in  the  sense  which  this  material  favors, 
if  the  larger  amount  of  material  is  also  scattered  through  a  larger 
space.  Chemical  changes  of  this  kind,  therefore,  are  not  accelerated 
by  increasing  the  mere  quantity  of  any  ingredient,  but  only  by  increas- 
ing the  concentration  of  its  molecules.  Hence,  in  practice,  we  find 
that  the  distribution  of  the  molecules  in  a  system  which  has  reached  a 
state  of  equilibrium  is  at  once  affected  if  we  alter  the  molecular  con- 
centration of  any  of  its  components.  Thus,  if,  in  the  action  of  hydro- 
gen upon  iodine,  we  introduce  into  ike  same  space  an  extra  amount 
of  hydrogen,  this  facilitates  the  formation  of  hydrogen  iodide  by  in- 
creasing the  possibilities  of  encounter  between  hydrogen  and  iodine, 
while  at  the  same  time  it  does  not  affect  (cf.  p.  88)  the  number  of 
encounters  in  a  given  time  between  hydrogen  iodide  molecules  which 
result  in  the  reverse  transformation.  The  proportion  of  hydrogen 
iodide  formed,  therefore,  from  a  given  amount  of  iodine  will  be 
greater,  although  the  potential  maximum  has  not  been  altered  since 
the  quantity  of  one  ingredient  only  has  been  increased.  The  intro- 
duction of  an  excess  of  iodine  would  have  had  precisely  the  same 
effect. 


250  INORGANIC   CHEMISTRY 

It  is  easy  to  illustrate  this  experimentally.  If  ferric  chloride 
and  ammonium  thiocyanate  are  mixed  in  aqueous  solution,  a  liquid 
containing  the  soluble,  blood-red  ferric  thiocyanate  is  produced; 
FeCl,  +  3NH4CNS^Fe(CNS)g  +  3NH4Cl.  The  action  is  a  rever- 
sible one.  Now,  if  the  two  just-named  salts  are  mixed  in  very  dilute 
solution  in  the  proportions  required  by  the  equation,  say  by  adding 
20  c.c.  of  a  deci-normal  solution  (p.  149)  of  each  to  several  liters  of 
water,  a  pale-reddish  solution  is  obtained.  When  this  is  divided 
into  four  parts,  and  one  is  kept  for  reference,  the  addition  of  a  little 
of  a  concentrated  solution  of  ferric  chloride  to  one  jar,  and  of  ammo- 
nium thiocyanate  to  another,  will  be  found  to  deepen  the  color  by 
producing  more  of  the  ferric  thiocyanate.  On  the  other  hand,  mix- 
ing a  few  drops  of  concentrated  ammonium  chloride  solution  with  the 
fourth  portion  will  be  found  to  remove  the  color  almost  entirely  on 
account  of  its  influence  in  accelerating  the  backward  change. 

We  may  state  (see  below)  as  a  general  principle  that,  in  reversible 
actions  between  substances  in  homogeneous  mixture,  the  amount  of  a 
chemical  change  taking  place  in  a  given  time  will  be  dependent  upon 
the  molecular  concentration  of  each  ingredient.  It  will  also  naturally 
depend  upon  the  intrinsic  affinity  (cf.  p.  111).  It  will  likewise  vary 
•with  the  temperature,  since  the  affinity  is  affected  by  change  in  tem- 
perature. 

The  phrase  "  active  mass  "  is  sometimes  employed  instead  of  the  words  "  molec- 
ular concentration."  It  is  somewhat  misleading,  however,  for,  as  we  have  seen, 
it  is  not  on  the  mass  of  a  substance,  but  on  the  quantity  of  it  in  a  given  volume, 
that  the  speed  of  the  action  depends. 

Law  Connecting  Molecular  Concentration  and  Speed  of  Re- 
action. —  These  principles  must  now  be  stated  in  somewhat  more  pre- 
cise terms.  We  confine  our  attention  to  homogeneous  mixtures,  in 
the  first  place,  and  to  non-reversible  actions. 

The  concentration  of  the  molecules  is  usually  expressed,  for  each 
substance,  in  terms  of  the  number  (whole  or  fractional)  of  moles 
(gram-molecular  weights,  p.  194)  of  the  substance  contained  in  a  liter 
of  the  whole  mixture.  There  is  the  same  number  of  molecules  in  a 
mole  of  every  substance,  namely,  the  number  of  molecules  in  32  g.  of 
oxygen  (cf.  p.  195).  Hence  the  number  of  moles  per  liter  defines  the 
concentration  of  the  substance  in  terms  of  this  number  of  molecules 
in  a  liter  as  the  unit  of  concentration. 

Thus,  in  a  solution  containing  25.4  g.  of  free  iodine  (^  of  a  formula 


CHEMICAL  EQUILIBRIUM  251 

weight,  I2)  per  liter,  the  solution  is  0.1  molar  (p.  149),  and  the  molec- 
ular concentration  of  the  iodine  is  0.1.  When  the  substance  is  a  gas, 
the  concentration  of  the  molecules  is  proportional  to  the  partial  press- 
ure of  the  gas.  Now,  one  mole  of  a  gas  occupies  22.4  liters  at  one 
atmosphere  pressure  (and  0°).  Hence,  when  one  mole  of  a  gas  is  con- 
tained in  1  liter,  and  its  molecular  concentration  is  therefore  1,  it  exer- 
cises 22.4  atmospheres  partial  pressure.  When  the  partial  pressure  of 
one  gas  in  a  mixture  is  two  atmospheres,  its  molecular  concentration  is 

,-fi  «  0-09- 

In  an  action  of  the  form  A  -f-  B  — >  C  -f  D,  let  c^  and  c2  represent 

the  number  of  moles  per  liter  of  the  materials  A  and  B,  respectively, 
at  any  stage  of  the  interaction.  Then  the  speed  (S)  of  the  forward 
action,  expressed  in  moles  of  A  and  B  transformed  per  unit  of  time 
(say,  per  hour),  is  found  to  be  defined  by  the  relation  ct  x  c2  x  F  —  S. 
Of  course  ct  and  c2  (and  therefore,  also,  S)  diminish  steadily,  for  the 
materials  A  and  B  are  being  progressively  used  up.  S,  therefore,  at 
any  moment  is  the  amount  which  would  be  transformed  if  the  concen- 
trations present  at  that  moment  were  artificially  maintained  during  the 
whole  hour.  .F  expresses  the  intrinsic  activity  (affinity)  propelling  the 
action,  which  is  independent  of  concentration.  If  unit  concentrations 
are  taken,  ct  —  c2  =  1,  and  therefore  F  =  S.  Thus  F,  the  activity,  is 
always  represented  numerically  by  the  speed  in  moles  per  unit  of  time 
when  the  concentration  of  each  ingredient  is  unity.  For  rapid  actions 
a  shorter  unit  of  time  than  the  hour  may  be  used. 

In  practice  the  value  of  F  for  any  action  can  easily  be  determined, 
because  cv  c2,  and  S  can  be  measured.  As  a  check  on  the  results, 
different  concentrations  will  be  used  at  the  same  temperature,  and  the 
amounts  transformed  in  measured  times  will  be  observed  in  each 
case.  We  can  then  Calculate  from  the  data  of  each  set  the  speed 
(moles  transformed  per  hour  or  minute)  which  would  be  shown  by 
constantly  maintained  unit  concentrations  of  the  materials.  The 
answers  are  the  values  of  F  based  upon  different  concentrations  of  the 
same  substances,  and  they  agree  closely  with  one  another.  The  values 
of  F  for  several  different  chemical  actions,  however,  may  differ  widely, 
and  are  measures  of  the  relative  activity  of  each. 

If  the  action  is  more  complex,  the  same  principles  apply.  Thus,  if  an  expres- 
sion for  the  speed  of  the  action  2A  +  SB  +  C  — *  D  +  E  is  required,  we  must  con- 
sider this  equivalent  to  A  +A  +  B  +  B+B+C-^>D  +  E.  Here  the  speed 
forwards  (S)  =  cx  x  q  x  c2  x  C2  x  C2  x  C3  x  F,  or  c*cfc3F  =  S. 

It  need  hardly  be  added  that,  clearly,  when  the  concentration  of  any  ingredient 


252  INORGANIC   CHEMISTRY 

becomes  zero  (say  by  all  of  it  entering  into  combination),  or  when  the  affinity  is 
zero,  the  speed  S  must  become  zero,  that  is  to  say,  no  action  takes  place. 

The  relation  of  the  theoretical  speed  with  constant  concentration, 
used  in  the  above  formulae,  to  that  which  is  observed  with  diminishing 
concentration  is  rather  complex.  For  an  action  of  the  form  A  — >  B  +  C, 
where  the  change  in  only  one  molecule  constitutes  the  action,  if  ^  is 
the  initial  molecular  concentration  of  A,  and  x  is  the  fraction  of  this 
which  is  transformed  in  the  time  t, 

s  =  log. -*--•-*. 


When  two  molecules  have  to  interact,  the  formula  is  still  more 
plex.  If  the  substances  are  present  in  equivalent  proportions,  their 
molecular  concentrations  in  this  special  case  are  alike,  and  may  each 
be  represented  by  cr  The  relation  is  then : 

*=_?_-,.*. 

ci  (ci  ~  x) 

The  mathematical  derivation  of  these  relations  will  be  found  in  any 
work  on  physical  chemistry. 

The  law  of  molecular  concentration  is,  therefore  :  In  every  chemi- 
cal experiment,  the  speed  of  the  action  at  any  moment  is  proportional 
to  the  first,  or  some  higher  power  of  the  molar  concentration,  for  the 
time  being,  of  each  interacting  substance,  and  to  the  affinity  at  work. 

The  following  illustration  (see  also  Sulphurous  acid)  will  make  all 
this  clearer.  When  arsine  AsH3  (q.v.)  is  heated  at  310°,  it  decomposes 
gradually  into  hydrogen  and  arsenic  : 

AsH3-+As  +  1JH2.* 

ci 

The  action  is  not  appreciably  reversible.  The  arsenic  assumes  the 
solid  form.  The  gas  is  inclosed  in  a  tube  which  is  kept  in  a  bath  at 
310°,  and  a  manometer  shows  changes  in  pressure.  Since,  as  the  action 
proceeds,  1J  molecules  of  hydrogen  take  the  place  of  each  molecule  of 
arsine,  the  total  pressure  slowly  increases.  Every  increase  of  1  mm.  in 
the  pressure  is  the  result,  therefore,  of  an  addition  of  3  mm.  partial 
pressure  of  hydrogen  and  a  reduction  of  2  mm.  in  the  partial  pressure 

*  Ordinarily  we  should  write  the  equation  2AsH3  — >  2As  -f  3H2.  But  this 
form  would  make  the  speed  proportional  to  Cj2,  and  calculation  then  gives  us  in- 
constant values  for  S.  The  reason  for  this  irregularity  is  given  in  works  on 
physical  chemistry. 


CHEMICAL    EQUILIBRIUM 


253 


of  the  arsine.  The  molecular  concentrations  are  proportional  to  the 
pressures,  and  change,  therefore,  in  the  same  ratios.  The  observations 
(first  two  columns),  together  with  the  data  deduced  from  the  first  two 
by  calculation,  were  as  follows  : 


MOLECULAB  CONCENTRATIONS. 

Total. 

AsH3  Transf  md. 

0 

784.84 

0.02159  (c.) 

3 

878.50 

0.02416 

6.00514 

0.0906 

4 

904.06 

• 

8 

987.19 

.... 

.     .     . 

.     .     . 

We  must  first  ascertain  the  molecular  concentration  of  the  arsine 
corresponding  to  the  observed  pressure  at  the  beginning.  We  remember 
that  at  22.4  atmospheres,  or  22.4  x  760  mm.  and  0°,  the  concentration 
of  a  gas  has  the  value  1  (p.  251).  The  actual  initial  pressure  784.84 


mm.  at  310°  would  become,  at 


>  or  367.5  mm.     The 


9  (310  +  273) 

Qf*rr   f» 

molecular  concentration  is  here,  therefore,  ^-^ — -?M\  or  0.02159  moles 

per  liter.  After  3  hours,  some  hydrogen  has  been  formed.  The 
pressure  has  increased  to  878.50  mm.  Reducing  as  before,  this  repre- 
sents a  molecular  concentration  of  all  ingredients  of  0.02416  moles  per 
liter.  The  increase  is  0.00257.  This,  as  was  demonstrated  above,  corre- 
sponds to  a  loss  of  2  x  0.00257,  or  0.00514  moles  per  liter  of  arsine. 
Now,  employing  the  formula  given  above,  we  find  the  speed  per  hour : 


This  result  means  that,  if  the  concentration  of  the  arsine  were  to  be 
maintained  at  the  initial  value  by  continual  renewal  of  the  waste, 
then  0.0906  (9.06  per  cent)  of  the  initial  amount  would  be  decomposed 
in  an  hour.  Using  the  pressures  at  4  and  at  8  hours,  the  reader  will 
obtain  by  calculation,  practically  the  same  value  for  S.  Other  experi- 
ments with  still  different  concentrations,  provided  the  temperature 
was  the  same  (p.  250),  would  likewise  give  the  same  result.  Hence, 
when  we  take  unit  concentration  (1  mole  per  liter),  c  =  1,  and  the  ex- 
pression c^F  =  S  becomes  F  =  0.0906.  Thus  the  affinity  or  activity 


254  INORGANIC   CHEMISTRY 

of  the  action  may  be  measured  with  any  concentration,  and  expressed 
in  moles  transformed  per  hour  with  unit  concentration.  Similar 
measurements  with  other  actions  then  enable  comparisons  of  their 
relative  activities  to  be  made  (see  Exercise  9,  end  of  this  chapter). 

The  Condition  for  Chemical  Equilibrium.  —  Let  us  now  take  a 
reversible  action, 

A<=±B+  C, 

^      c2       c,, 

in  which  cv  c2,  cs  are  the  molecular  concentrations  after  the  condition 
of  equilibrium,  has  been  established.  The  above  mathematical  form  of 
statement  for  the  influence  of  concentration  and  affinity  on  speed  of 
transformation,  being  true  for  all  concentrations,  will  be  true  for  this 
set.  The  speed  of  the  forward  and  reverse  actions  will  be, 

Sl  =  clFl    and    S2  =  c2c8F2, 

respectively,  where  F^  is  the  intrinsic  tendency  of  A  to  decompose, 
and  Fz  the  tendency  of  B  and  C  to  combine.  As  the  concentrations 
have  so  adjusted  themselves  that  an  equal  amount  of  material  is  being 
transformed  each  way,  we  have 

^-cfyF,   or  J  =  ^. 

•*  2  °1 

F 

-=f- ,  being  the  ratio  of  two  constants,  is  constant  (  =  k).     This  is  the 

ratio  of  the  affinities  driving  the  opposed  actions,  and  is  known  as  the 
affinity  constant  of  the  reversible  chemical  change.  If,  for  example, 
its  value  is  -J-,  then  the  speeds  of  the  two  actions,  if  each  were  to  pro- 
ceed unimpeded  (say  in  separate  vessels)  with  constantly  maintained 
unit  concentrations  of  the  materials,  would  be  in  the  ratio  F^ :  F2 
or  1  :  4.  From  this  it  will  be  seen  that  measurement  of  the  concentra- 
tions present  in  a  system  which  has  reached  equilibrium  gives  us  the 
data  for  calculating  the  value  of  this  ratio.  In  other  words,  it  gives  us 
the  means  of  ascertaining  the  relative  magnitudes  of  the  intrinsic 
affinities  of  the  opposed  actions. 

We  may  apply  this  to  the  data  given  (p.  237)  for  hydrogen  iodide.     The  equa- 
tion is 

2HI<=±H2  +  I2, 
c,         c2     c3, 

and  Cj,  c2,  and  c3  are  the  molecular  concentrations  after  the  system  has  come  to  rest. 
The  speed  of  the  forward  action  (SJ  is  cJ2Fl  and  that  of  the  reverse  action  (S2)  is 
cac3F2.  As  concentrations  are  now  such  that  the  speeds  are  equal,  we  have 


CHEMICAL  EQUILIBRIUM  255 


. 

*2  Cl 

Now,  at  448°,  with  equivalent  quantities  of  the  two  elements,  nearly  0.8  (more 
exactly,  79  per  cent)  of  the  weight  of  each  is  in  the  5orm  HI,  and  0.2  in  the  mix- 
ture H2  +  I2.  Thus  in  every  100  molecules,  80  are  KI(Ci),  10  are  H2(c2),  and  10 
are  I2(c3).  Thus  k  =  O.I2  -s-  0.82  =  ^¥.  That  is  to  say,  the  union  of  hydrogen  and 
iodine  would  take  place  with  64  times  as  great  a  speed  as  the  dissociation  of  hy- 
drogen iodide  if  each  action  could  proceed  without  reversal  and  under  identical 
conditions.  Or,  in  terms  of  the  kinetic  theory,  the  collisions  of  the  H2  and  I2 
molecules  result  many  times  more  often  in  chemical  change  than  do  collisions 
of  HI  molecules. 

The  case  of  hydrogen  iodide  is  comparatively  simple  because  the  volume  is  not 
altered  by  the  progress  of  the  action  (see  below).  The  expansion  when  phos- 
phorus pentachloride  (p.  213)  dissociates  compels  us  to  take  account  of  the 
volume.  The  equation  is  : 

PC16  +±  PC13  +  C12,  and  ^  =  k  =  ^  . 
GI  c2         ca  F2  cx 

If  one  gram  molecule  of  the  substance  is  taken,  and  x  is  the  proportion  dissociated, 

i  £  j  _  ,p 

and  v  the  volume  occupied  by  the  whole,  then  c2  =  c3  =  -  and  Cj  =  --    Thus 

k  =  —  ---  —  —  •     Now  at  250°  (and  760  mm.),  for  example,  0.8  of  the  whole  weight 

of  material  is  dissociated  :  x  =  0.8,  1  —  x  =  0.2.  Hence  k  =  0.82  -r-  0.2t?  =  3.2  -=-  v. 
To  obtain  the  value  of  v  we  note  that  a  gram  molecule  at  760  mm.  and  0°  occupies 
22.4  liters.  At  260°  it  occupies  22.4  x  (250  +  273)  -r-  273  1.  But  this  mass  of  gas 
contains  0.8  more  molecules  because  of  dissociation,  and  its  volume  is,  therefore, 

1.8  x  22.4(250+  273)-^  273  .=  0  =  77.21.    Thus  k  =  3.2  -r-  77.2  =  ~    Otherwise 

stated,  25  Fl  =  Fr  That  is  to  say,  the  union  of  the  trichloride  and  chlorine  would 
proceed  twenty-five  times  as  fast  as  the  dissociation,  if  each  of  the  three  sub- 
stances was  present  in  unit  concentration,  and  each  action  could  proceed  inde- 
pendently without  reversal. 

The  Effect  of  Changes  of  Volume  on  Chemical  Equilibrium. 

—  Our  applications  of  the  theory  of  equilibrium  will  be  chiefly  to  dis- 
solved bodies,  and  hence  the  effect  on  the  equilibrium  point  of  changes 
in  volume  (by  dilution  or  the  reverse)  will  require  frequent  considera- 
tion. Now  dilution,  for  example,  diminishes  opportunities  for  en- 
counters between  the  substances  on  both  sides  of  the  equation.  In  the 
first  of  the  above  illustrations,  increasing  the  volume  decreases  the 
rate  at  which  the  chlorine  and  the  trichloride  can  combine.  Since, 
however,  the  speed  of  the  dissociation  depends  on  the  state  of  the 
PC16  molecules  only,  and  is  unaffected  by  their  nearness  to  or  remote- 
ness from  one  another,  the  forward  action  will  not  be  weakened  at  all. 


256  INORGANIC   CHEMISTRY 

Hence,  dilution  increases  the  degree  of  dissociation.  In  general, 
change  in  volume  will  affect  the  equilibrium  point  whenever  there  are 
more  molecules  on  one  side  of  the  equation  than  on  the  other. 

In  mathematical  terms,  when  we  change  the  volume  to  -  times  its 
former  value  (n  whole  or  fractional),  the  concentration  changes  n 
times.  The  equation  for  equilibrium  then  becomes,  momentarily, 
ncz  x  nc8  -5-  TiCj  ^  k,  or  nc2cs  -f-  ct  ^  k.  To  restore  the  value  of  the 
expression  to  equality  with  k,  change  must  occur  in  the  concentrations 
C2,  cs,  and  c1.  When  n  is  <  1,  that  is,  when  the  volume  increases, 
some  PC15  must  pass  into  the  form  PC1S  and  C12  until  c2'c3'  -*-  c/=  k, 
as  before.  That  is,  dilution  increases  the  degree  of  dissociation. 

In  the  case  of  hydrogen  iodide,  and  in  all  others  where  the  number  of  molecules 
taking  part  in  the  direct  and  reverse  actions  is  the  same,  change  in  the  volume  of 
the  system  has  no  effect  on  the  position  of  the  equilibrium  point.  Thus  dilution 
diminishes  the  chance  of  encounter  between  two  HI  molecules  to  the  same  extent 
that  it  interferes  with  encounters  between  H2  and  I2  molecules.  Conversely, 
increase  in  all  concentrations,  by  diminution  of  volume,  favors  both  actions  equally. 
Hence,  at  448°,  79  per  cent  of  HI  will  always  be  present  at  last,  whatever  the 
volume  occupied  by  a  given  amount  of  the  materials.  In  mathematical  terms,  if  we 
diminish  the  volume  n  times  (n  whole  or  fractional),  we  increase  the  concentration 
of  each  constituent  n  times.  The  values  become  ncn  nc2,  and  nc3  respectively, 

*--• 


Heterogeneous  Equilibrium.  —  A  modification  of  the  above 
conceptions  is  necessary  when  the  mixture  is  not  homogeneous.  If, 
for  example,  one  of  the  constituents  is  present  as  a  solid  or  a  gas,  in 
greater  amount  than  can  be  dissolved  by  the  liquid  in  which  alone  the 
chemical  change  takes  place,  then,  according  to  the  definition  of 
saturated  solution  (p.  158),  the  concentration  of  the  dissolved  material 
will  be  constant  at  a  given  temperature  as  long  as  physical  equilibrium 
between  the  solid  and  the  solution  is  maintained.  This  is  a  case 
especially  likely  to  occur  when  slightly  soluble  (so-called  "  insoluble") 
bodies  (of.  p.  146)  are  concerned. 

The  same  reasoning  applies  also  to  very  slightly  volatile  solids. 
The  concentration  of  the  vapor  of  a  solid  body  present  in  excess  (meas- 
ured by  its  vapor  pressure)  will  be  constant  so  long  as  the  temperature 
is  fixed,  and  interaction  with  a  superincumbent  gas  will  take  place 
chiefly  through  the  vapor. 

In  both  these  cases  the  concentrations  of  the  active  parts  of  the 


CHEMICAL    EQUILIBRIUM  257 

slightly  soluble  and  slightly  volatile  bodies,  respectively,  are  not  sub- 
ject to  variation  —  they  are  constant.     Thus,  with  the  action, 


2Ba02^±2BaO  +  02, 


the  concentrations  of  the  vapor  of  BaO2(<?1)  and  of  BaO(c2)  are  constant, 
and  that  of  oxygen  (cs)  alone  is  subject  to  alteration.  We  have,  there- 
fore, 

^  =  k,  or  cs  =  £  k, 

Cl  C2 

in  which,  since  el9  c2,  and  k  are  constant,  c8  must  be  constant  also.  But 
the  pressure  of  a  gas  is  proportional  to  its  molecular  concentration, 
according  to  Avogadro's  hypothesis.  Therefore,  in  this  action,  the 
pressure  of  the  oxygen  (the  dissociation  pressure)  should  be  constant 
irrespective  of  the  extent  to  which  the  dissociation  has  progressed. 
Observation  shows  that  this  is  the  case.  When  barium  dioxide  is 
maintained  at  a  definite  temperature,  the  pressure  of  the  oxygen  rises 
to.  a  fixed  value  and  remains  constant.  Any  pressure  of  oxygen  above 
the  fixed  value  forces  the  gas  into  combination  until  no  barium  mon- 
oxide is  left  ;  any  pressure  below  this  value  permits  the  dioxide  to 
decompose  until  none  remains.  This  explains  the  Brin  method  of 
making  oxygen  (p.  63),  in  which  a  temperature  is  chosen  at  which 
the  dissociation  pressure  is  approximately  equal  to  that  of  the  partial 
pressure  of  oxygen  in  the  air  (£  atmosphere).  All  hydrates,  as  we 
have  seen  (pp.  121-123),  behave  in  a  precisely  similar  way,  and  fur- 
nish numberless  confirmations  of  this  application  of  the  law  of  molecu- 
lar concentration. 

Applications  in  Chemistry  :  Displacement  of  Equilibria  by 
Changes  Affecting  Concentration.  —  Of  special  interest  to  the 
chemist  are  the  conditions  under  which  the  equilibrium  point  may  be 
displaced  and  more  nearly  complete  realization  of  one  of  the  two 
opposed  changes  may  be  brought  about. 

We  have  just  seen  (p.  249)  that  one  way  in  which  a  reversible  action 
may  be  forced  nearer  to  completion  in  one  direction  or  the  other  is  the 
introduction  of  an  excess  of  one  of  the  ingredients  contributing  to  the 
action.  This  method  of  displacing  the  equilibrium  point,  however, 
cannot  be  very  effective  unless  it  is  possible  to  introduce  an  exceedingly 
large  excess  of  the  selected  ingredient  in  a  high  degree  of  molecular 
concentration,  since  this  operation  does  not  in  any  way  effect  or,  in 


258  INORGANIC    CHEMISTRY 

particular,  restrain  the  reverse  action  which  is  continually  undoing  the 
work  of  the  forward  one.  A  much  more  effective  means  of  furthering 
such  actions  is  found  in  removing  one  of  the  components  of  the  system,. 
Any  agency  which  could  remove  the  free  iodine  as  fast  as  it  was 
formed  in  the  decomposition  of  hydrogen  iodide,  for  example,  would 
entirely  stop  the  reproduction  of  the  compound  and  so  would  enable 
the  dissociation  (2HI  +±  H2  -f-  I2)  to  run  to  completion. 

This  might  be  realized  *  by  causing  one  end  of  a  sealed  tube  charged 
with  hydrogen  and  iodine,  after  the  contents  had  settled  down  to  a 
condition  of  equilibrium,  to  project  from  the  bath  in  which  the  whole 
had  been  kept  at  448°  (Fig.  72,  which  is  simply  diagrammatic).  By 

cooling  this  end,  a  large  part  of  the 

^/~~  /\          21    per   cent   of   free  iodine   would 

quickly  be  condensed  in  it  to  the  solid 
form,  while  the  hydrogen  would  re- 
main gaseous.  Only  the  trace  of 
vapor  which  cold  iodine  gives  would 


FlGt  72-  then  be  available  to  interact  with  the 

hydrogen   and  reproduce    hydrogen 

iodide.  Meanwhile  the  decomposition  of  the  latter  would  go  on,  and 
thus,  eventually,  almost  all  the  iodine  would  be  found  free  in  one  end 
of  the  tube,  and  the  hydrogen,  all  free  likewise,  would  occupy  the  rest. 
By  this  purely  mechanical  adjustment  the  chemical  change  would  in 
this  way  be  carried  from  21  per  cent  completion  to  almost  absolute 
completion. 

If,  on  the  other  hand,  arrangements  were  made  to  have  powdered 
marble,  in  a  sealed  bulb  of  thin  glass,  inclosed  in  the  tube,  we  might 
imagine  the  very  opposite  effect  of  the  above  to  be  produced.  The 
breaking  of  the  bulb  of  marble,  when  equilibrium  had  been  reached, 
would  provide  means  for  the  removal  of  all  the  hydrogen  iodide  ,t  while 
the  hydrogen  and  iodine  would  still  be  gaseous.  Thus,  the  com- 
pound having  been  removed,  there  would  be  no  reverse  action  to 
compensate  for  the  union  of  the  elements.  The  whole  material  would, 
therefore,  soon  have  passed  through  the  form  HI.  Hence,  by  another 

*  For  another  illustration,  see  under  Ammonia. 

t  The  hydrogen  iodide  would  be  destroyed  by  interaction  with  the  marble: 

2HI  +  CaCO3  — >  CaI2  +  C02  +  H2O. 

The  calcium  iodide  is  a  solid.  The  two  gases,  carbon  dioxide  and  water  vapor, 
do  not  interact  with  hydrogen  or  with  iodine,  and  would  not,  therefore,  inter- 
fere with  the  formation  of  fresh  hydrogen  iodide. 


CHEMICAL   EQUILIBRIUM  259 

mechanical  arrangement,  an  action  which  ordinarily  could  progress  to 
only  79  per  cent  would  be  turned  into  a  complete  one. 

In  every-day  chemical  work,  since  our  object  is  usually  to  prepare 
some  one  substance,  chemists  either  avoid  chemical  changes  which  are 
notably  reversible,  or  adjust  the  conditions  so  that  the  reverse  of  the 
action  which  they  desire  is  prevented.  In  consequence  of  this,  when 
carrying  out  the  directions  for  making  familiar  preparations,  the  fact 
that  such  actions  are  reversible  at  all  very  readily  escapes  our  notice. 
Arranging  the  conditions  so  that  the  separation  of  a  solid  body  by  pre- 
cipitation, or  the  liberation  of  a  gas,  takes  place,  are  the  two  com- 
monest ways  of  rendering  a  reversible  action  complete.  Excellent 
examples  of  both  of  these  are  furnished  by  the  chemical  change  used 
in  producing  hydrogen  chloride  by  the  interaction  of  salt  and  sul- 
phuric acid  (p.  179),  the  discussion  of  which  should  once  more  be 
studied  attentively. 

The  escape  of  one  member  of  a  system  engaged  in  chemical  interaction, 
because  it  is  gaseous  or  solid,  and  in  either  case  immiscible  with  the  rest  of  the 
members  of  the  system,  is  the  commonest  cause  of  the  obstruction  of  one  direc- 
tion of  a  reversible  action  and  the  triumph  of  the  other.  This,  as  we  have  seen, 
is  the  combined  result  of  the  natural  behavior  of  a  system  in  chemical  equilibrium, 
and  of  the  physical  properties,  particularly  the  solubility,  of  the  members  of  the 
system.  Two  rules,  attributed  to  Berthollet,  have  been  made,  however,  to  describe 
these  special  cases  of  a  broader  principle.  Unfortunately,  it  is  difficult  so  to  word 
them  that  they  shall  be  entirely  unambiguous  and  entirely  correct. 

The  "rule of  precipitation,"  for  example,  might  read:  When  certain  classes  of 
materials  are  brought  together  in  solution,  if  an  exchange  of  radicals  would  produce 
an  insoluble  body,  this  exchange  will  occur.  But  then  the  fact  is  that,  in  such 
cases,  the  exchange  always  occurs  to  some  extent  whether  any  product  is  insoluble 
or  not.  The  insolubility  is  responsible  only  for  the  greater  completeness  of  the 
exchange.  Crude  statements  to  the  effect  that  "  when  an  insoluble  body  can 
be  formed,  it  will  be  formed,"  when  close  scrutiny  shows  them  to  possess  any 
definite  meaning  whatever,  are  grossly  misleading.  They  suggest  that  insolubility 
is  a  sort  of  especially  desirable  career  on  which  the  elements  are  ambitious  of 
entering. 

All  forms  of  these  so-called  laws  are  objectionable,  because  they  necessarily 
suggest  that  the  positive  direction  of  the  action  is  assisted  by  the  immiscibility  of 
the  product,  and  this  is  the  precise  converse  of  the  fact.  The  immiscibility  does 
nothing  at  all  towards  assisting  the  formation  of  the  insoluble  substance  itself,  but 
does  whatever  it  can  towards  preventing  the  destruction  of  that  substance,  once  it 
is  formed,  by  hampering  the  negative  action. 

Affinity  vs.  Solubility.  —  The  question  of  the  relation  of  affinity 
to  the  apparently  much  greater  efficiency  of  one  of  the  directions  of 
some  reversible  actions,  may  now  be  put  in  a  much  clearer  light  (pp. 


260  INORGANIC  CHEMISTRY 

181,  110,  and  this  Chap.).  The  whole  of  the  possibilities  of  progress 
for  any  action  are  expressed  by  a  function  (p.  251)  of  the  form  c^F  =  S. 
If  any  one  of  the  variables,  say  one  of  the  concentrations  (Cj),  is  neg- 
ligible, the  product  must  be  small,  irrespective  of  the  values  of  the 
other  factors.  Thus  the  feebleness  of  a  chemical  action  only  shows 
that  the  product  of  all  the  variables  is  minute,  and  not  that  the 
affinity  factor  per  se  is  of  small  magnitude. 

Displacement  of  Equilibria  by  Changes  in  Temperature: 
Van't  Hoff's  Law  and  Le  Chatelier's  Law.  —  Since  the  affinities 
driving  the  two  opposed  actions  are  differently  affected  by  change  in 
temperature,  the  value  of  their  ratio  alters,  and  the  equilibrium  point, 
which  depends  on  this  ratio,  suffers  displacement.  It  is  found  that  the 
direction  of  the  displacement  depends  on  which  of  the  two  actions 
absorbs  heat,  and  which  gives  it  out.  In  a  system  which  is  in  equilib- 
rium, of  the  two  opposed  interactions,  that  one  which  is  endothermal 
is  promoted,  while  that  which  is  exothermal  is  resisted,  by  raising  the 
temperature,  and  vice  versa.  This  is  van't  Hoff  s  law  of  mobile  equilib- 
rium. Thus  the  dissociation  of  phosphorus  pentachloride  (p.  208)  is 
endothermal  (p.  27),  and  hence  becomes  greater  at  higher  tempera- 
tures. All  thermal  dissociations  are  of  this  kind.  Again,  the  inter- 
action of  steam  and  iron  (p.  110)  is  exothermal,  and  so  the  higher  the 
temperature,  the  more  conspicuous  the  opposite  action  becomes.  In 
view  of  this  generalization,  we  can  understand  the  formation  of  endo- 
thermal substances  of  an  unstable  character  at  high  temperatures. 
Thus,  ozone  (q.v.)  is  produced  by  electrical  discharges,  and  cyanogen 
(q.v.)  is  formed  in  the  blast-furnace. 

The  principle  applies  also  to  physical  equilibria.  Thus,  as  the 
temperature  rises,  a  compound  which  gives  out  heat  in  dissolving  is 
less  soluble  in  a  solution  already  almost  saturated  with  the  compound; 
while  one  which  absorbs  heat  in  dissolving  is  more  soluble  in  such  a 
solution.  For  example,  anhydrous  sodium  sulphate  gives  out  heat 
in  dissolving  (p.  165),  and  its  solubility  diminishes  (p.  158)  with  ris- 
ing temperature,  while  with  hydrated  sodium  sulphate  just  the  con- 
verse is  observed.  So,  also,  the  vaporization  of  a  liquid  absorbs  heat, 
and  hence  the  concentration  of  vapor  it  can  maintain  increases  with 
rise  in  temperature. 

The  above  law  is  a  particular  case  of  a  still  more  general  one  frequently 
called  Le  Chatelier's  law :  If  some  stress  (for  example,  by  change  of  tem- 
perature, pressure,  or  concentration)  is  brought  to  bear  on  a  system  in  equi- 


CHEMICAL  EQUILIBRIUM  261 

librium,  by  which  the  equilibrium  is  displaced,  the  equilibrium  is  dis- 
placed in  that  direction  which  tends  to  undo  the  effect  of  the  stress. 

Thus,  pressure  causes  ice  to  melt  because  the  water  which  is  formed  occupies  a 
smaller  volume,  and  the  change  therefore  tends  to  relieve  the  pressure.  So,  also, 
the  production  (say,  by  a  chemical  interaction)  of  a  large  amount  of  a  body  in 
solution,  by  which  its  concentration  is  increased  beyond  saturation,  leads  to  crys- 
tallization (or  precipitation)  of  the  excess. 

Exercises.  — 1.  What  is  the  molecular  concentration  of  the  oxygen 
in  the  air  (pp.  155,  250),  of  the  nitrogen  in  the  air,  of  the  aqueous  vapor 
above  water  at  10°  and  at  20°  (p.  116),  of  a  solution  containing  one 
formula-weight  of  sodium  chloride  in  10  liters,  of  a  solution  containing 
65  g.  of  hydrogen  iodide  in  250  c.c.  ? 

2.  What  are  the  partial  pressures   of   the   three    components  of 
phosphorus  pentachloride  vapor  at  250°  and  760  mm.  (p.  207)  ?    What 
are  their  molecular  concentrations  ? 

3.  Using   the   model   on   p.   254,  study  the   dissociation  of  KIg 
(p.  235),  of  iodine  vapor  (p.  236),  and  of  hydrogen  iodide  (p.  237), 
and   the  formation  of   ferric  thiocyanate   (p.  250).      Show   in    each 
case  the  effect  on  the  system  of  increase  in  volume  without  change  in 
the  amount  of  material  (p.  255). 

4.  Using  the  model  on  p.  257,  study  the  dissociation  of  mercuric 
oxide  (p.  12),  assuming  the  compound  to  be  involatile,  and  the  inter- 
action of  iron  and  steam  (p.  110).     Why  can  magnetic  oxide  of  iron 
be -reduced  completely  by  a  stream  of  hydrogen  (p.  110),  and  iron 
oxidized  completely  by  a  current  of  steam  (p.  99)  ? 

5.  Wliat  actions  in  Chap,  xiv  are  complete  for  the  same  reason 
that  the  action  of  sulphuric  acid  on  salt  (pp.  178-180)  is  so  ? 

6.  Why  is  the  formation  of  the  following  substances  complete : 
silver  chloride  (p.  186),  and  hydrogen  chloride  and  water  by  union 
of  the  elements  ? 

7.  What  inference  should  you  draw  from  the  fact  that  the  solu- 
bilities  of   potassium   nitrate,   sodium   chloride,  and    Glauber's   salt 
(p.  157)  increase  with  rise  in  temperature  (p.  260),  and  from  the  fact 
that  those  of  calcium  citrate  (p.  156)  and  triethylamine  decrease  with 
rise  in  temperature  ? 

8.  Is  the  heat  of  solution  of  lead  nitrate  (p.  157)  positive  or  nega- 
tive? 

9.  Carry  out  the  calculation  of  $  for  4  and  8  hours  (p.  253). 


262  INORGANIC   CHEMISTRY 

SUMMARY  OF  PRINCIPLES. 

The  summary  of  some  of  the  chief  principles  of  the  science  (p.  188)  may  now 
receive  several  important  additions.  For  the  sake  of  completeness,  reference  to 
the  periodic  system  is  made  in  No.  21,  to  isomers  in  No.  22,  and  to  the  phase  rule 
in  No.  23,  although  these  subjects  have  not  yet  been  taken  up.  As  before,  all 
hypothetical  matters  have  been  excluded  as  far  as  possible. 

15.  That  weight  of  each  substance  which  in  the  gaseous  condition  occupies  the 
same  volume  as  32  grams  of  oxygen,  temperature  and  pressure  being  alike  for 
both  (namely,  22.4  liters  at  0°  and  760  mm.),  is  taken  as  the  chemical  unit  of 
weight  for  the  substance,  and  is  known  as  its  molar  iceight  (p.  195). 

16.  That  weight  of  each  element  which  is  the  greatest  common  measure  of  the 
quantities  of  the  element  found  in  the  molar  weights  of  its  compounds  is  taken 
as  the  chemical  unit  of  weight  for  the  element,  and  is  known  as  its  atomic  weight. 
This  weight  has  the  property  described  in  6  (p.  188). 

The  composition  of  each  substance  is  expressed  in  terms  of  the  atomic  weights 
as  units,  and  the  sum  of  the  atomic  weights  is  multiplied  by  an  integer,  when 
necessary,  so  as  to  equal  the  molar  weight  (p.  204). 

17.  The  number  of  equivalent  weights  of  hydrogen  which  combine  with,  or  are 
displaced  by  the  atomic  weight  of  an  element  is  called  the  valence  of  the  element 
(p.  103). 

18.  The  speed  of  every  interaction  is  a  function  of  the  first,  or  some  higher 
power  of  the  molar  concentration  of  each  interacting  substance  (p.  262). 

19.  Substances  undergoing,  at  a  fixed  temperature,  an  interaction  which  is 
reversible,  reach  a  condition  of  equilibrium.     The  final  proportions  of  the  mate- 
rials are  such  that  the  speeds  (see  18)  of  the  opposed  actions  are  equal  (p.  254). 

20.  Van't  Hoff's  law  and  Le  Chatelier's  law  (p.  260). 

21.  Each  element  has  its  own  set  of  chemical  relations  (pp.  177,  226)  :  e.g.,  it 
can  exist  in  combination  with  certain  other  elements ;  it  has  a  certain  valence, 
and  may  have  more  than  one  valence  ;  it  confers  certain  properties  on  its  com- 
pounds as  a  class  ;  it  resembles  certain  other  elements  in  several  of  these  respects 
(e.gr.,  the  halogens),  and  differs  from  others,  in  a  way  more  or  less  definitely 
described  by  its  place  in  the  periodic  system  (g.t).). 

In  complex  cases,  the  inter-relations  of  the  elementary  units  in  a  compound, 
and  the  relations  of  the  compound  to  other  compounds  (see  No.  22),  are  represented 
graphically  by  formulae  based  upon  an  hypothesis  of  molecular  structure  (p.  224). 

22.  Identical  combinations  of  matter  may  constitute  more  than  one  compound 
substance  (isomers,  see  Urea).     These  may  have  equal  molar  weights  (optical 
and  structural  isomers),  or  they   may  have  different  molar  weights  (pp.    204, 
242). 

23.  In  a  system  in  equilibrium  the  number  of  components  plus  two  equals  the 
number  of  phases  plus  the  number  of  degrees  of  freedom  (Phase  rule, 


CHAPTER   XVI 
OXIDES  AND  OXYGEN  ACIDS  OP  THE  HALOGENS 

THE  chief  subjects  of  practical  importance  touched  upon  in  this 
chapter  are  connected  with  bleaching  powder  (CaCl(OCl)),  and  potas- 
sium chlorate  (KC103)  and  perchlorate  (KC104).  Hence  our  attention 
will  be  largely  directed  to  the  modes  of  making  these  substances  and  to 
their  relations  to  one  another.  Incidentally,  we  shall  encounter  many 
actions  of  a  complex  and,  to  us,  more  or  less  novel  kind. 

Compounds  of  Chlorine  Containing  Oxygen.  —  The  following 
are  the  names  and  formulae  of  the  substances  : 

HC1O  Hypochlorous  acid,  C12O  Hypochlorous  anhydride, 

(HC102)  Chlorous  acid,  

C1O2  Chlorine  dioxide, 

HC1O3' Chloric  acid, 

HC104  Perchloric  acid,  C12O7  Perchloric  anhydride. 

There  are  also  compounds  of  metals  with  the  negative  radicals  of 
these  acids.  Of  this  nature  are  the  three  substances  mentioned  in  the 
first  paragraph.  Chlorous  acid  is  itself  unknown,  but  potassium  chlorite 
(KC102)  and  some  other  derivatives  have  been  made. 

The  two  anhydrides  (p.  71 ),  when  brought  into  contact  with  water, 
combine  with  it  to  form  the  acids  opposite  which  they  stand  in  the 
table.  Chlorine  dioxide  (q.v.),  however,  is  not  related  to  any  one  acid 
in  this  way. 

All  these  compounds  differ  from  most  that  we  have  hitherto  dis- 
cussed inasmuch  as  not  one  of  them  can  be  made  by  direct  union  of 
the  simple  substances. 

Nomenclature  of  Acids  and  Salts-  —  When  several  compounds 
closely  related  in  composition,  like  the  above  acids,  are  known,  a  sys- 
tematic method  of  naming  them  is  used.  The  terminations  -ous  and 
-ic  indicate  smaller  and  larger  proportions  of  oxygen  respectively. 
For  compounds  below  or  above  those  two  in  their  degree  of  oxidation, 
the  prefixes  hypo-  and  per-  are  employed. 

263 


264  INORGANIC   CHEMISTRY 

When  the  radicals  (p.  93)  contained  in  the  acids  are  combined 
with  metals,  the  compounds  are  spoken  of  as  salts  of  the  respective 
acids.  Thus,  KC103  is  described  as  the  potassium  salt  of  chloric  acid. 
The  specific  names  for  these  salts  are  distinguished  by  terminations 
corresponding  to  those  of  the  acids  : 

KC1O  Potassium  hypochlorite,  KC103   Potassium    chlorate, 

KC102  Potassium  chlorite,  KC1O4  Potassium  perchlorate. 

The  termination  -ite  corresponds  to  -ous,  -ate  to  -ic.  This  principle  is 
applied  systematically,  so  that  the  salts  of  sulphuric  and  sulphurous 
acids,  for  example,  are  called  sulphates  and  sulphites  respectively. 

Compounds  containing  no  oxygen  receive  the  termination  -ide. 
Thus,  KC1  is  potassium  chloride,  FeS  is  ferrous  sulphide. 

Salts  and  Double  Decomposition.  —  We  have  just  been  using 
the  word  salt  in  a  general  sense.  It  is  the  class  name  for  a  set  of  sub- 
stances which  includes  common  salt  or  sodium  chloride  (NaCl),  potas- 
sium nitrate  (KN08),  sodium  sulphate  (Na^SC^),  silver  chloride  (AgCl), 
potassium  chlorate  (KC108),  etc.  The  majority  of  the  substances  used 
in  elementary  chemistry  belong  to  this  class.  They  receive  the  name 
because  in  certain  important  chemical  respects  they  behave  like  com- 
mon salt.  For  example,  when  sodium  chloride  is  treated  with  acids, 
such  as  sulphuric  acid  (p.  178)  or  phosphoric  acid  (p.  179),  hydrogen 
chloride  is  liberated.  Actions  of  this  kind  consist  in  an  exchange  of 
radicals  and  are  all  reversible.  An  action  of  the  same  type,  although 
outwardly  different,  is  that  of  sodium  chloride  and  silver  nitrate  in 
aqueous  solution  (p.  99).  Here  we  have  two  salts  interacting  instead 
of  an  acid  and  a  salt,  and  we  get  a  precipitate  instead  of  a  gas.  But 
the  interchange  of  radicals  is  exactly  similar. 

Now  salts  in  general  behave  in  these  respects  in  the  same  way  as 
does  common  salt.  They  interact  with  acids  or  other  salts,  particu- 
larly in  solution,  in  such  a  way  that  an  exchange  of  radicals  takes 
place.  In  the  first  case,  a  salt  and  an  acid,  and  in  the  second  case 
two  salts,  are  produced.  These  actions  are  all  reversible.  Acids  differ 
thus  from  salts  only  in  the  fact  that  one  of  their  radicals  is  hydrogen. 
Hence  they  are  frequently  called  hydrogen  chloride,  hydrogen  sulphate 
(H2S04),  and  so  forth.  It  may  be  added  that  bases  (p.  119),  like  potas- 
sium hydroxide  (KOH),  interact  reversibly  with  salts  and  acids, 
exchanging  radicals  after  the  same  fashion. 

All  salts  are  named  according  to  the  radicals  which  they  contain. 


OXIDES   AND   OXYGEN  ACIDS  OF  THE    HALOGENS  265 

Thus,  all  containing  —  S04  are  sulphates.  Conversely,  when  the  name 
of  a  salt  is  given,  the  formula  can  be  written  down  at  once.  In  doing 
this,  however,  regard  must  be  had  to  the  valence  of  the  radicals 
(p.  104). 

In  view  of  the  reversibility  of  most  of  the  interactions  of  salts, 
acids,  and  bases,  we  encounter  completed  changes  chiefly  when  pre- 
cipitation occurs,  or  when  one  product  is  volatile  (p.  259).  If  neither 
of  the  products  formed  by  the  exchange  of  radicals  is  insoluble,  the 
reversibility  of  the  action  prevents  our  obtaining  anything  but  a 
mixture.  Only  those  double  decompositions  which  involve  more  or  less 
insoluble  or  volatile  substances  are  thus  of  use  for  preparing  salts.  The 
action  of  sodium  chloride  on  silver  nitrate  (p.  99)  is  an  example.  The  sil- 
ver chloride  is  almost  completely  insoluble,  while  the  sodium  nitrate  pro- 
duced by  the  change  remains  dissolved.  By  nitration  we  obtain  the 
silver  chloride  as  a  powder,  while  the  evaporation  of  the  filtrate  gives 
us  the  soluble  product.  This  sort  of  action  can  be  used,  therefore, 
either  for  the  preparation  of  a  soluble  or  an  insoluble  substance.  If 
the  problem  is  to  make  a  soluble  product,  then  we  must  arrange  an 
action  between  two  substances,  each  containing  one  of  the  two  required 
radicals,  and  possessing  two  other  radicals,  which,  when  united,  give 
an  insoluble  body.  This  plan  is  illustrated  frequently  in  what  follows. 

Hypochlorites.  —  Since  none  of  the  acids  in  our  list  can  be  made 
directly  from  their  elements,  we  generally  have  to  prepare,  first,  the 
corresponding  salt.  From  the  salt,  by  double  decomposition,  the  acid 
is  then  secured.  Hence,  in  each  case,  the  salts  will  be  discussed 
first. 

Chlorine  interacts  slightly  with  water  (p.  176),  producing  small 
quantities  of  hydrogen  chloride  and  hypochlorous  acid  (equation  (1), 
below).  The  action  is  very  strongly  reversible.  That  is  to  say,  since 
the  last  two  substances  interact  very  vigorously  to  reproduce  chlorine 
and  water,  the  direct  action  does  not  make  much  progress. 

When,  however,  some  substance  which  can  interact  with  these 
products  is  added  to  the  solution  of  chlorine,  or  when  chlorine  gas  is 
simply  passed  into  an  aqueous  solution  of  such  a  substance,  displace- 
ment of  the  equilibrium  point  at  once  occurs  (p.  258).  Now  potas- 
sium hydroxide  is  a  suitable  substance.  It  interacts  almost  completely 
in  solution  with  both  the  products  of  this  action,  producing  potassium 
chloride  (2)  and  potassium  hypochlorite  (3),  according  to  the  last  two 
of  the  following  equations  : 


266  INORGANIC   CHEMISTRY 

CL,  +  H20        ±+  HC1  +  HOC1,  (1) 

HC1  +  KOH    _>KC1  +  H20,  (2) 

HOC1  +  KOH  -»  KOC1  +  H/).*  (3) 

Thus,  omitting  the  water  which  appears  both  among  products  and 
initial  substances,  and  the  two  acids  which  are  used  up  as  quickly  as 
they  are  produced  by  the  first  action,  we  get,  by  addition  of  the  three 
partial  equations  (cf.  p.  228),  the  final  equation  : 

CL,  +  2KOH  ->  KOI  +  KOC1  +  H20. 

This  sort  of  action  does  not  give  a  pure  hypochlorite,  but  for  some 
purposes  the  presence  of  the  chloride  in  the  solution  is  not  objection- 
able. 

Bleaching  powder,  CaCl(OCl),  is  manufactured  on  a  large  scale  by 
an  action  exactly  like  the  above.  The  neutralization  of  a  molecule  of 
each  of  the  two  acids,  however,  can  be  accomplished  by  a  single  molecule 
of  slaked  lime  (calcium  hydroxide,  Ca(OH)2),  since  the  latter  contains 
two  hydroxyl  (OH)  groups.  The  hydroxide  can  be  applied  either  in 
the  dry  form  or  mixed  with  some  water  as  a  paste.  The  separate 
actions  and  final  equation  are  as  follows  : 

CLj    +  H20      ±^HC1    +      HOC1  (1) 

/OH  +  HC1  /Cl 

CaNQH  ^Ca\  +  2H'0 


CLj+Ca(OH)2-^CaCl(OCl)  +  H20 

Bleaching  powder  (see  Calcium)  is  a  salt  of  calcium  involving  two 
different  acids  (a  mixed  salt).  This  condition,  again,  does  not  inter- 
fere with  the  application  of  the  substance  commercially.  A  method 
of  obtaining  pure  hypochlorites,  however,  will  be  found  below. 

Hypochlorites  change  into  chlorates  (q.v.)  when  heated.  They 
may  also  give  off  oxygen,  2CaCl(OCl)  ->  2CaCLj  +  02.  Although 
this  decomposition  is  slow  in  cold  solutions  of  hypochlorites,  or  when 
they  are  preserved  in  the  dry  form,  it  may  be  hastened  by  means  of 
catalytic  agents.  The  addition  of  a  little  cobalt  hydroxide  (q.v.) 
to  bleaching  powder  solution  causes  rapid  evolution  of  oxygen.  The 
hypochlorites  are  manufactured  because  hypochlorous  acid,  which  is 
used  in  bleaching,  can  readily  be  made  from  them.  The  acid  itself 
will  not  keep,  except  when  largely  diluted,  and  consequently  cannot 
be  transported  conveniently. 

*  The  interaction  of  potassium  hydroxide,  or  any  other  base,  with  any  acid  to 
produce  a  salt  and  water,  is  called  neutralization  (cf.  p.  186). 


OXIDES  AND   OXYGEN   ACIDS   OF  THE   HALOGENS  267 

Preparation  of  Hypochlorous  Acid:  Hypochlorous  Anhy- 
dride. —  1.  The  common  method  of  obtaining  the  acid  is  by  double 
decomposition,  using  some  other  acid  (p.  264).  Even  from  such  a 
mixture,  or  mixed  salt,  as  is  produced  by  the  action  of  chlorine  on 
a  base  (p.  266),  the  acid  may  be  obtained  in  fairly  pure  condition. 
Thus,  with  nitric  acid  (an  active  acid),  we  have,  simultaneously,  the 
two  reversible  actions : 

(  KOC1  +  HN08  <=»  KN03  +  HOC1, 
\  KC1  +  HN08     <=±  KN03  +  HC1. 

But  hypochlorous  acid  is  a  feeble  acid,  while  hydrochloric  acid  is  an 
active  one,  so  that  in  the  former  action  the  reversing  tendency  is  very 
slight,  while  in  the  latter  it  is  vigorous.  Hence,  by  adding  nitric  acid, 
in  amount  barely  sufficient  for  the  liberation  of  the  hypochlorous  acid 
alone,  and  doing  this  in  a  very  dilute  solution,  the  object  is  attained. 
The  potassium  chloride  is  hardly  affected.  By  gently  warming  the 
liquid  a  dilute  solution  of  hypochlorous  acid  can  be  distilled  off.  The 
operation  may  be  performed  even  more  successfully  by  use  of  a  weak 
acid,  instead  of  an  active  one  like  nitric  acid.  Boric  acid  (q.v.)  is 
suited  to  the  purpose. 

2.  Advantage  may  be  taken  of  the  feebleness  of  hypochlorous  acid 
in  another  way.  When  powdered  chalk  (CaC03)  is  added  to  chlorine 
water,  or  chlorine  is  passed  into  water  holding  chalk  in  suspension,  only 
the  hydrogen  chloride  has  any  appreciable  action  upon  the  chalk.  It 
gives,  by  exchange  of  radicals,  calcium  chloride  and  carbonic  acid 
(equation  (2),  below).  The  latter  decomposes  immediately  with  effer- 
vescence, yielding  water  and  carbon  dioxide  gas  (equation  (3)). 

In  this  particular  example  of  equation-making  (cf.  p.  229)  the 
three  partial  equations,  in  their  simplest  forms,  may  not  be  added 
together.  Before  this  can  be  done  we  must  always  see  that  the  terms 
which  appear  on  both  sides  of  the  equations,  and  are  not  actual 
products,  will  disappear  by  cancellation.  Here,  the  2HC1  in  equation 
(2)  will  not  cancel  with  the  first  product  in  equation  (1),  unless  the 
whole  of  this  equation  is  multiplied  by  two.  The  multiplication  has 
therefore  been  effected : 

2CL,  +  2H20     £5  2HC1+  2HOC1  (1) 

2HC1  +  CaC08  fc?  CaCL,  +  H2C08  (2) 

H2C08<^H2O  +  C02 (3) 

2CL,  +  CaC08  +  H80  -» CaCL,  +  CO,  \  +  2HOC1. 


268  INORGANIC   CHEMISTRY 

This  action  gives  a  solution  containing  calcium  chloride  and  hypo- 
chlorous  acid,  and,  by  distillation,  as  in  the  first  method,  a  dilute  solu- 
tion of  the  acid  may  be  secured. 

3.  The  anhydride  of  hypochlorous  acid  (C120)  may  be  obtained  by 
passing  chlorine  gas  over  mercuric  oxide.  For  this  purpose  precipi- 
tated mercuric  oxide  (q-v.)  must  be  used,  and  it  should  be  heated  in 
advance  in  order  to  coarsen  the  grain  of  its  particles  and  so  diminish 
the  speed  with  which  it  is  acted  upon.  Each  of  the  constituents  of  the 
oxide  combines  with  chlorine  : 

HgO  +  20^  ->  HgCL,  +  01,0. 

The  mercuric  chloride  then  unites  with  another  formula-weight  of  the 
mercuric  oxide  to  form  a  compound  HgO,  HgCl2,  which  remains  in 
the  tube.  The  chlorine  monoxide  is  a  reddish-yellow  gas.  It  may 
be  reduced  to  the  liquid  form,  and  boils  at  +  5°.  Both  the  gaseous 
and  liquefied  forms  of  it,  the  former  when  heated,  the  latter  when 
touched  by  paper  or  dust,  decompose  into  the  constituents  with  explo- 
sion. The  gas  dissolves  in  water  very  easily  (200:  1,  by  vol.).  The 
yellow  solution  of  hypochlorous  acid  which  results, 


has  a  strong  odor  of  chlorine  monoxide.   The  combination  is  reversible, 
and,  when  the  liquid  is  warm,  a  little  of  the  gas  escapes. 

Properties  of  Hypochlorous  Acid.  —  1.  Hypochlorous  acid 
cannot  be  made,  excepting  in  solution,  or  kept,  excepting  in  dilute 
solution.  This  is  in  consequence  of  its  tendency  to  decompose  in  three 
different  ways,  one  of  which  has  just  been  mentioned  (see  3  and  4 
below). 

2.  As  an  acid  it  neutralizes  (p.  266)  active  bases,  giving   hypo- 
chlorites. 

3.  If  the  solution  is  concentrated,  much  of  the  hypochlorous  acid 
changes   gradually   into   chloric   acid   and   hydrogen   chloride.     This 
occurs  even  in  the  dark. 

3HC10-»HC108+2HCL 

4.  When  the  solution  is  warmed,  but  more  especially  when  it  is 
exposed  to  sunlight,  oxygen  is  evolved  rapidly. 

2HOC1->2HC1  +  0. 


OXIDES   AND  OXYGEN  ACIDS  OF  THE   HALOGENS  269 

This  decomposition  always  takes  place  whether  the  acid  is  present 
alone  in  the  water,  or  along  with  other  substances.  Hence,  the  solu- 
tion of  chlorine  in  water,  which  contains  a  small  amount  of  hypo- 
chlorous  acid,  on  being  exposed  to  the  bright  sunlight  gives  off  bubbles 
of  oxygen  in  rapid  succession.  This  decomposition,  since  it  re- 
moves one  of  the  interacting  substances  in  the  reverse  action,  C12  + 
H20  ±=;  HC1  -f-  HOC1,  enables  the  interaction  of  chlorine  and  water  to 
go  on  to  completion.  Consequently,  the  final  liquid  contains  nothing 
but  hydrochloric  acid  and  water.  Leaving  out  the  intermediate  steps 
again,  the  action  appears,  therefore,  to  be  simply  a  decomposition  of 
water  by  chlorine. 

2CL,  +  2H20-+  4HC1  +  O2. 

5.  In  consequence  of  the  ease  with  which  it  gives  up  oxygen,  hypo- 
chlorous  acid  is  a  strong  oxidizing  agent. 

Hypochlorous  Acid  as  an  Oxidizing  Agent :  Bleaching. —  Both 
iodine  and  bromine  are  oxidized  by  hypochlorous  acid,  the  former 
much  more  rapidly  than  the  latter,  2HOC1  -f  I2  — » 2HOI  +  CL,. 
Further  oxidation  to  HI03  occurs  immediately.  Although  iodine  has 
less  affinity  for  hydrogen  than  has  chlorine  (p.  239),  this  action  shows 
that  the  relation  towards  oxygen  is  just  the  opposite.  Here  the  iodine 
goes  into  combination  and  the  chlorine  is  displaced. 

It  is  on  account  of  its  oxidizing  power  that  hypochlorous  acid  is 
used  commercially  in  bleaching.  It  is  not  applied  to  paints,  which 
are  chiefly  mineral  substances,  but  to  complex  compounds  of  carbon, 
such  as  constitute  the  coloring  matters  of  plants  and  of  those  artificial 
dyes  whose  manufacture  has  now  become  so  gigantic  an  industry.  It 
should  be  understood  that  the  great  majority  of  the  complex  com- 
pounds of  carbon  are  colorless.  Even  a  slight  chemical  change,  affect- 
ing only  one  or  two  of  the  atoms  in  a  complex  molecule,  is  thus  almost 
sure  to  give  a  colorless  or  much  less  strongly  colored  material.  Indigo 
(C16H101S'202),  which  has  a  deep-blue  color,  is  an  example  of  a  vegetable 
dye  which  is  also  made  artificially.  Hypochlorous  acid  oxidizes  it  to 
isatin,  a  yellow  substance  relatively  pale  in  color : 

C16H10N202  +  2HOC1  ->  2C8H6N02  +  2HCL 

In  ways  just  as  definite  as  this,  hypochlorous  acid  will  change  the 
composition  of  other  colored  substances,  although,  since  we  do  not 
know  the  formulae  of  all  these  substances,  we  cannot  always  write 


270  INORGANIC   CHEMISTRY 

equations  for  the  actions.  Thus,  the  interaction  by  which  chloro- 
phyll, the  green  coloring  matter  of  plants,  is  bleached  is  doubtless 
similar  to  the  above,  although  the  formulae  of  materials  concerned  are 
unknown. 

On  account  of  the  hypochlorous  acid  which  is  already  present  in 
chlorine  water,  this  solution  is  a  very  efficient  bleaching  agent.  The 
removal  of  this  one  of  the  factors  in  the  reverse  action  (p.  265)  enables 
more  of  the  acids  to  be  produced  from  the  chlorine  and  water  until  the 
whole  of  the  former  has  been  consumed. 

As  a  rule,  bleaching  is  actually  carried  out  by  liberating  hypochlo- 
rous acid  from  bleaching  powder  by  means  of  sulphuric  acid : 

,OC1     EL  HOC1 

Ca  +        S04-»CaS04+ 

XC1       H/  HC1 

Of  course,  temporarily,  most  of  the  hypochlorous  acid  interacts  with 
the  hydrochloric  acid  to  give  chlorine  and  water,  but,  as  the  residual 
hypochlorous  acid  loses  its  oxygen,  the  secondary  action  is  again  dis- 
placed backwards  until  the  chlorine  is  all  used  up. 

The  yarn  or  cloth  is  first  cleansed  from  fatty  or  oily  material  by 
boiling  with  soap  solution.  It  is  then  immersed  in  bleaching  powder 
solution,  and  finally  in  dilute  sulphuric  acid.  Both  solutions  must  be 
very  weak  in  order  that  no  interaction  may  occur  with  the  fabric  it- 
self. The  last  two  processes  may  be  repeated,  if  the  brownish  or 
yellowish  coloring  material  has  not  disappeared  after  the  first  treat- 
ment. 

Hypochlorous  acid  can  be  used  to  bleach  linen  or  cotton,  because 
the  body  of  these  materials,  apart  from  the  small  amount  of  coloring 
matter,  is  composed  of  compounds  containing  nothing  but  carbon, 
hydrogen,  and  oxygen.  These  compounds  are  very  slowly  affected  by 
hypochlorous  acid,  unless  too  strong  a  solution  is  used,  or  the  exposure 
to  its  influence  is  too  long.  That  chemical  interaction  does  occur  is 
shown  by  the  "rotting"  of  goods  which  have  not  been  washed 
thoroughly  after  bleaching.  Wool,  silk,  and  feathers,  on  the  other 
hand,  are  composed  largely  of  compounds  containing  nitrogen  in  addi- 
tion to  the  above  three  elements.  Their  constituent  material  interacts 
as  easily  with  hypochlorous  acid  as  do  the  traces  of  coloring  substances. 
Hence,  since  the  fabric  itself  would  be  attacked  by  this  agent,  different 
means  of  bleaching  have  to  be  used  for  materials  of  this  class. 

It  should  be  understood  that  a  cold  dilute  solution  of  hypochlorous 


OXIDES  AND   OXYGEN   ACIDS   OF  THE    HALOGENS  271 

acid  may  be  kept  almost  indefinitely  and  will  not  give  up  its  oxygen 
spontaneously.  The  transfer  takes  place  when,  and  only  when,  the 
acid  comes  in  contact  with  some  substance  capable  of  uniting  with 
oxygen. 

Thermochemistry  of  Hypochlorous  Acid.  —  As  we  have  seen 
(p.  27),  chemical  changes  which  proceed 'spontaneously  to  completion 
are  accompanied  by  a  transformation  of  chemical  energy  into  some 
other  form  of  energy.  Hence,  a  substance,  or  system  of  substances, 
which  undergoes  such  a  change,  possesses  more  chemical  energy 
and  activity  before  the  change  than  after  it.  In  consequence,  if 
some  given  chemical  change  uses  the  products  of  such  an  action, 
and  can  be  brought  about  by  the  employment  of  the  original  sub- 
stance, the  employment  of  the  latter  will  involve  a  greater  liberation 
of  energy,  and  will  therefore  be  more  likely  to  secure  the  consumma- 
tion of  the  change  in  question. 

The  decomposition  of  Jiypochlorous  acid  and  of  chlorine  monoxide 
are  cases  where  there  is  a  very  marked  difference  between  the  amount 
of  chemical  energy  in  the  original  substances  and  in  the  products  of 
decomposition,  hydrogen  chloride  and  free  oxygen  in  the  first  case,  and 
free  chlorine  and  oxygen  in  the  second.  Hence  the  changes  into  these 
substances  sometimes  are  of  the  nature  described  as  explosive.  A  more 
important  fact,  however,  is  that,  on  this  account,  hypochlorous  acid 
and  chlorine  monoxide  are  more  active  oxidizing  agents  than  is  free 
oxygen  gas.  The  energy  liberated  in  the  decomposition  of  the  hypo- 
chlorous  acid  has  to  be  added  (p.  78)  to  that  which  free  oxygen 
could  give,  if  performing  the  same  oxidation,  in  order  that  the  total 
fall  in  energy,  which  measures  the  tendency  of  the  action  to  take 
place,  may  be  estimated.  Hence,  substances  that  are  not  affected  by 
free  oxygen  may  be  changed  instantly  by  hypochlorous  acid.  This 
explains,  for  example,  the  oxidation  by  hypochlorous  acid  of  many 
carbon  compounds,  including  those  which  are  colored,  when  atmos- 
pheric air  is  without  action.  Thus,  the  heat  liberated  in  the  oxidation 
of  indigo  to  isatin  by  oxygen  gas,  if  it  could  be  carried  out,  would  be 
1800  cal.  The  much  greater  heat  liberated  when  hypochlorous  acid 
is  used,  we  obtain  by  adding  the  thermochemical  equations : 

2HC10  =  2HC1  +  20  +  18,600  cal. 
C16H10lSr202+  20  =  2C8H6N02    +    1800  cal. 
C16H10N302  -f  2HC10  =  2C8H6NOa  +  2HC1  +  20,400  cal. 


272  INORGANIC   CHEMISTRY 

The  following  thermochemical  equations  give  a  rough  Idea  of  the  relative 
oxidizing  powers  of  the  chief  oxygen  acids  of  the  halogens  : 

HC10,  Aq  -  HC1,  Aq  +    O  -f     9300  cal.,   or  +  9,300 


HC10,,  Aq  =  HC1,  Aq  +  3O  -f  15,300  cal.,  or  -f  5,100 
HC1O4,  Aq  =  HC1,  Aq  +  4O  +  700  cal.,  or  +  170 
HBrO,,  Aq  =  HBr,  Aq  +  3O  +  15,000  cal.,  or  +  5,000 
HIO,,  Aq  =  HI,  Aq  -f  30  -  42,900  cal.,  or -14, 300 


cal.  for 
each  atomic 
''weight  of 
oxygen*. 


HIO4,    Aq  =  HI,    Aq  +  4O  -  34,500  cal.,   or-  8,600  , 

Formerly  a  different  explanation  for  actions  like  that  of  hypochlorous  acid, 
when  it  behaves  as  an  oxidizing  agent,  was  offered.  It  was  suggested  that  the 
oxygen  was  first  liberated  from  the  acid  (HOC1— »HC1  +  O),  and  that  the  single 
atoms  of  the  element  so  produced  were  more  active  than  molecular  oxygen.  This 
oxygen,  which  was  supposed  to  interact  in  the  moment  of  its  production,  was 
called  nascent  oxygen.  But  it  will  be  seen  that  such  an  explanation  is  entirely 
unnecessary.  The  activity  of  the  hypochlorous  acid  on  account  of  its  large  store  of 
free  energy  sufficiently  accounts  for  the  facts  (see  Nascent  hydrogen) . 

Simultaneous,  Independent  Chemical  Changes  in  the  Same 
Substance.  —  As  we  have  seen,  hypochlorous  acid  undergoes  three 
different  changes.  Some  molecules  decompose  into  water  and  chlorine 
monoxide  (p.  268),  while  others  give  chloric  acid  and  hydrogen  chlo- 
ride, and  still  others  hydrogen  chloride  and  oxygen.  Since  the 
same  molecule  cannot  undergo  more  than  one  of  these  different 
changes,  it  follows  that  the  actions  are  independent  of  one  another. 
This  is  shown  by  the  fact  that  in  sunlight  the  third  predominates, 
while  in  the  dark  it  falls  far  behind  the  second.  Since  the  relative 
quantities  of  the  products  vary,  the  several  simultaneous  actions 
cannot  be  put  in  the  same  equation.  The  fundamental  property  of 
an  equation  is  to  show  the  constant  proportions  by  weight  between 
every  pair  of  substances  in  it.  Hence  three  separate  equations  are  re- 
quired in  the  present,  and  in  all  similar  cases  where  all  the  proportions 
are  not  constant  (cf.  p.  231  and  see  Per  chlorates).  Successive  actions, 
like  (1)  preceded  by  (2)  in  the  next  section  (cf.  p.  266),  however,  may 
be  combined  in  one  equation,  since  in  them  all  the  proportions  must 
necessarily  be  constant.  These  equations  are  interlocked,  for  (1) 
consumes  what  (2)  produces. 

Chlorates.  —  Like  hypochlorous  acid  itself,  the  hypochlorites  turn 
into  chlorates.  Thus,  when  chlorine  is  passed  into  a  warm,  concen- 
trated solution  of  potassium  hydroxide,  and  particularly  when  an 
excess  of  chlorine  is  used,  the  hypochlorite  changes  into  chlorate  as 
fast  as  it  forms  : 

3KC10  ->  RC108  +  2KC1.  (1) 


OXIDES    AND   OXYGEN   ACIDS   OF   THE    HALOGENS  273 

To  secure  the  three  molecules  of  the  hypochlorite,  the  equation  form- 
erly given  (p.  266)  must  be  tripled : 

3C12  +  6KOH  ->  3KC1  +  3KC1O  +  3H2O.  (2) 

When  these  are  added,  and  the  intermediate  substance  is  left  out,  the 
final  equation  is  obtained : 

3C12  -h  6KOH  ->  KC108  +  5KC1  +  3H2O. 

When  the  solution  is  cooled,  the  chlorate  crystallizes  out. 

This  action  involves  converting  five-sixths  of  the  valuable  potas- 
sium hydroxide  into  the  relatively  less  valuable  potassium  chloride. 
Hence,  in  practice,  the  makers  carry  out  the  corresponding  action 
with  calcium  hydroxide.  They  then  add  potassium  chloride  to  the 
resulting  solution,  containing  calcium  chloride  and  calcium  chlorate 
(Ca(C103)2).  The  potassium  chlorate,  formed  by  double  decomposi- 
tion, crystallizes  when  the  solution  is  cooled. 

All  chlorates  are  at  least  moderately  soluble  in  water.  Potas- 
sium chlorate  is  used  in  making  fireworks,  explosives,  and  matches. 
An  intimate  mixture  with  sugar  (C^H^C^)  burns  with  semi-explosive 
violence,  the  oxygen  of  the  salt  combining  with  the  carbon  and  hydro- 
gen to  form  carbon  dioxide  and  water.  Detonating  fuses  for  artillery 
are  made  of  a  mixture  of  this  salt  with  antimony  trisulphide  (y.v.). 

The  Separation  of  Substances  by  their  Solubility.  —  When 
neither  of  the  products  of  an  action  approaches  absolute  insolubility,  a  separation 
may  nevertheless  be  effected  more  or  less  perfectly  by  taking  advantage  of  differ- 
ence in  solubility.  Thus,  in  the  practical  method  of  making  potassium  chlorate, 
the  calcium  chloride  is  exceedingly  soluble,  while  the  potassium  chlorate  is  only 
moderately  so.  Then,  too,  the  solubility  of  the  latter  decreases  rapidly  as  the 
temperature  is  lowered  (Fig.  61,  p.  157).  Hence,  it  is  found  that  when  the 
mixture  is  cooled  to  — 18°  only  about  13.5  g.  of  potassium  chlorate  remain  dis- 
solved in  each  liter,  and  are  lost.  At  0°  the  loss  would  be  greater,  for  at  this 
temperature  a  liter  of  pure  water  would  kold  33.3  g.,  and  a  liter  of  this  solution 
would  contain  more  than  this  on  account  of  the  uncompleted  reversible  action 
(c/.  p.  264) : 

Ca  (C10,)2  +  2KC1  <=»  CaCl2  +  2KC1O,. 

It  will  be  seen  that  we  reason  as  if  the  solubility  of  each  substance  was  inde- 
pendent of  the  presence  of  other  dissolved  bodies  (p.  153). 

By  the  use  of  this  principle,  and  the  data  in  regard  to  solubility  in  Fig.  61 
(p.  157),  a  rough  idea  may  be  obtained  of  what  may  be  expected  in  any  given 
case.  From  the  diagram  the  solubilities  at  any  given  temperature  may  be  read. 
Suppose,  for  example,  the  question  is  in  regard  to  the  quantity  of  potassium 


274 


INORGANIC   CHEMISTRY 


chlorate  we  may  expect  to  obtain  from  3  g.  of  potassium  hydroxide  dissolved  in 
7  g.  of  water  (a  30  per  cent  solution).  From  the  equation  : 

3C12  +  6KOH  — »  KC1O3  +  5KC1  +  3H2O 
6X56  122.5         5X74.5 

we  find  that  336  g.  of  potassium  hydroxide  give  122.5  g.  of  chlorate  and  372.5  g.  of 
chloride.  Hence,  by  proportion,  3  g.  will  give  about  1  g.  and  3  g.  respectively. 
The  solubility,  read  from  the  diagram,  is  the  amount  of  the  salt  dissolved  by  100 
c.c.  of  water,  for  example,  56.5  g.  of  potassium  chloride  at  100°.  Some  of  the 
results  are  given  in  the  form  of  a  table  : 


POTASSIUM 
CHLOBIDE. 

POTASSIUM 
CHIX)KATB. 

Amount  form 
Solubility  at 
100°  in 
Solubility  at 
20°  in 
Solubility  at 
0°in 

ed  from  3  g.  KOH.       .     .     . 
100  c  c   Aq  .     .          ... 

3.0 
66.5 
4.0 
34.7 
2.6 
28.0 
2.0 

1.0 
56.5 
4.0 
7.5 
0.5 
3.3 
0.25 

*"      7  c  c.  Aq  

100  c  c.  Aq  ....     o     . 

100  c  c.  Aq  

7  c.c.  Aq  

Thus,  at  20°,  at  least  2.5  g.  of  the  3  g.  of  potassium  chloride  will  remain  dissolved, 
while  half  of  the  potassium  chlorate  will  crystallize  out.  If  the  solubilities  are 
examined,  it  will  be  seen  that  the  potassium  chlorate  is  even  more  easily  obtain- 
able in  pure  condition  when  calcium  chloride  takes  the  place  of  potassium 
chloride. 

Chloric  Acid.  —  This  acid  may  be  obtained  in  solution  in  water, 
by  adding  the  calculated  amount  of  diluted  sulphuric  acid  to  a  solution 
of  barium  chlorate  : 


Ba(C103) 


2HC10 


The  barium   sulphate,  being  insoluble,  is  removed  by  nitration  (cf. 
p.  265). 

The  solution  may  be  concentrated  (to  about  40  per  cent)  by  evapora- 
tion, but  must  not  be  heated  above  40°,  as  the  acid  decomposes  near  this 
temperature.  The  resulting  thick,  colorless  liquid  has  powerful  oxidiz- 
ing qualities,  setting  fire  to  paper  (made  of  cellulose,  C6H10O5)  which 
has  been  dipped  into  it.  It  converts  iodine  into  iodic  acid,  2HC103  -j- 
I2  —  -»  2HI08  -f  CLj.  When  warmed  beyond  40°  the  acid  decomposes, 
giving  chlorine  dioxide  and  perchloric  acid  (see  below). 

Making  of  Equations  Once  More.  —  The  equation  for  the  last 
action,  although  far  from  simple,  may  be  made  readily  by  use  of  a  device 


OXIDES   AND   OXYGEN   ACIDS   OF   THE   HALOGENS  275 

which  can  always  be  applied  where  an  oxygen  acid  gives  an  oxide.  The 
formula  of  the  initial  substance,  chloric  acid,  may  be  written  thus, 
HgOjCLjOg,  so  as  to  show  the  anhydride  (in  this  case  imaginary)  to 
which  it  is  related.  Now  the  products  are  C102  and  perchloric  acid, 
and  the  latter  may  be  written  H20)Clf)r  Disregarding  the  elements 
of  water,  we  perceive  that  some  CLj05  becomes  C1207,  while  the  rest  of 
the  CLjOg  furnishes  the  oxygen  for  this  change  and  itself  falls  to 
2C102.  Evidently  one  molecule  undergoing  the  former  change  will 
require  two  undergoing  the  latter  in  order  that  it  may  secure  the  two 
units  of  oxygen  : 

2H20,C1205  -4  2H20  +  4C102  (+  20),  (1) 

H20,C1205  (+  20)  -*  H20,C1207.  (2) 

Adding  and  dividing  by  2,  we  have  : 

3HC1O3  -»  H20  +  2C102  +  HC1O4. 

Chlorine  Dioxide :  Chlorous  Acid.  —  Chlorine  dioxide  is  a 
yellow  gas  which  may  be  liquefied,  and  boils  at  +  10°.  The  gas  and 
liquid  are  violently  explosive,  the  substance  being  resolved  into  its 
elements  with  liberation  of  much  heat.  It  is  formed  whenever  chloric 
acid  is  set  free,  and  hence  it  is  seen  when  a  little  powdered  potassium 
chlorate  is  touched  with  a  drop  of  concentrated  sulphuric  acid.  Con- 
centrated hydrochloric  acid  turns  yellow  from  the  same  cause  when 
any  chlorate  is  added  to  it.  These  actions  are  used  as  tests  for  chlo- 
rates, and  distinguish  them  from  perchlorates  (q.v.).  With  water,  chlo- 
rine dioxide  gives  a  mixture  of  chlorous  and  chloric  acids,  and  with 
bases  a  mixture  of  the  chlorite  and  chlorate. 

Per 'chlorates 9  Perchloric   Acid,  and  Perchloric  Anhydride. 

—  When  heated,  chloric  acid  and  chlorates  give  perchloric  acid  (p.  274) 
and  perchlorates  respectively.  The  chlorates  also  give  oxygen  at  the 
same  time  (p.  64)  : 

f2KC103  ->  2KC1  +  302, 

\4KC103  ->  3KC104  +  KC1. 

These  actions,  like  the  three  decompositions  of  hypochlorous  acid, 
are  independent,  and  proceed  simultaneously  (p.  272).  Their  relative 
speed,  however,  varies  with  the  temperature,  and  the  decomposition 
into  chloride  and  oxygen  may  completely  outrun  the  other  when  a 


276  INORGANIC   CHEMISTRY 

catalytic  agent  like  manganese  dioxide  is  added  (p.  65).  When  purt 
potassium  chlorate  is  heated  cautiously,  about  one-fifth  of  it  has  lost 
all  its  oxygen  by  the  time  the  rest  has  turned  into  perchlorate.  The 
mixture  may  be  separated  by  grinding  with  the  minimum  quantity  of 
water  which  will  dissolve  the  chloride  it  contains.  The  perchlorate, 
having  at  15°  less  than  one-twentieth  of  the  solubility  of  the  chloride, 
will  remain,  for  the  most  part,  undissolved.  The  perchloratee  are 
much  more  stable  (p.  119)  than  the  chlorates,  or  hypochlorites  :  they 
are  all  soluble  in  water,  and  they  are  used  in  making  matches  and  fire- 
works. 

Pure  perchloric  acid  explodes  when  heated  above  92°.  But,  like 
other  liquids,  its  boiling-point  is  lower  when  its  vapor  is  under  reduced 
pressure  (p.  118).  At  56  mm.  pressure  it  boils  at  39°,  a  temperature 
at  which  hardly  any  decomposition  is  noticeable.  Hence  the  acid  may 
be  made  by  mixing  potassium  perchlorate  and  concentrated  sulphuric 
acid  and  distilling  the  mixture  cautiously  in  a  vacuum  : 

KC1O4  +  H2S04  <=»  KHS04  +  HC10J. 

To  secure  the  requisite  low  pressure,  the  ordinary  distilling  apparatus 
(Fig.  16,  p.  38)  is  made  completely  air-tight,  and  is  connected  by  a 
branch  tube  with  a  water-pump. 

Perchloric  acid  is  a  colorless  liquid,  which  decomposes,  and  often 
explodes  spontaneously,  when  kept.  A  70  per  cent  solution  in  water 
is  perfectly  stable,  however.  Although  it  is  an  active  oxidizing  agent, 
it  is  not  so  active  as  chloric  acid,  and  does  not  oxidize  hydrogen  chlo- 
ride in  cold  aqueous  solution.  When  liberated  by  concentrated 
sulphuric  acid  it  does  not  at  once  give  the  yellow  chlorine  dioxide 
(p.  275). 

The  anhydride  (C1207)  may  be  prepared  by  adding  phosphoric  anhydride  to 
perchloric  acid  in  a  vessel  immersed  in  a  freezing  mixture,  P2O6  -f  2HC1O4  — >  2HPO3 
-f  C13O7.  Phosphoric  anhydride  is  often  used  in  this  way  for  removing  the 
elements  of  water  from  compounds.  By  gently  warming  the  mixture,  the  perchlo- 
ric anhydride  can  be  distilled  off.  It  is  a  colorless  liquid  boiling  at  82°  (760  mm.), 
and  exploding  when  struck  or  too  strongly  heated. 

Oxygen  Acids  of  Bromine.  —  No  oxides  of  bromine  have  been 
made,  but  the  acids  HBrO  (hypobromous  acid)  and  HBr03  (bromic 
acid)  and  their  salts  are  familiar. 

By  the  action  of  bromine  on  dilute,  cold  potassium  hydroxide 
solution,  the  bromide  and  hypobromite  are  formed  : 

Br2  +  2KOH  -» KBr  +  KBrO  +  H2O. 


OXIDES  AND  OXYGEN  ACIDS   OF  THE   HALOGENS  277 

When  the  solution  is  heated,  the  hypobromite  turns  into  bromate  and 
bromide.  The  actions  are  exact  parallels  of  the  corresponding  ones 
for  chlorine. 

Aqueous  bromic  acid  may  be  made  in  the  same  way  as  chloric  acid 
(p.  274),  or  by  the  action  of  chlorine  and  water  on  bromine  : 

6HC10  +  Br2  +  H2O->2HBr03  +  5HCL 

The  solution  is  colorless  and  has  powerful  oxidizing  properties.  Thus, 
it  converts  iodine  into  iodic  acid. 

2HBr08  +  I2  ->  2HI03  +  Br2. 

It  appears,  therefore,  that  iodine  has  more  affinity  for  oxygen  than  has 
bromine. 

Oxides  and  Oxygen  Acids  of  Iodine.  —  The  following  are  the 
acids  and  their  corresponding  salts  : 

[HIO    Hypoiodous  acid],  [KIO         Potassium  hypoiodite], 

HI08    Iodic  acid,  KIOS          Potassium  iodate, 

[HI04  Periodic  acid],  NaI04        Sodium  periodate, 

H6I06   Periodic  acid,  Na2H3I06  Disodium  periodate. 

The  substances  in  parenthesis  have  not  been  isolated.  There  is  one 
oxide,  I20B. 

lodates  and  Iodic  Acid.  —  The  potassium  and  sodium  salts  of 
iodic  acid  are  found  in  Chili  saltpeter.  They  may  be  made,  in  much 
the  same  fashion  as  are  the  chlorates  and  bromates  (p.  272),  by  adding 
powdered  iodine  to  a  hot  solution  of  potassium  or  sodium  hydroxide. 
There  is  evidence  that  hypoiodites  are  formed  in  cold  solutions,  but 
they  change  quickly  to  iodates. 

Iodic  Acid  is  formed  by  passing  chlorine  through  powdered 
iodine  suspended  in  water.  The  action  is  parallel  to  that  of  chlorine 
on  bromine  water.  A  still  better  way  is  to  boil  iodine  with  aqueous 
nitric  acid  (q.v.).  The  latter  gives  up  oxygen  readily,  and  is  here 
used  solely  on  this  account.  Hence  it  may  be  omitted  from  the 
equation,  only  the  oxygen,  of  which  it  is  the  source,  appearing : 

I2  +  H,0  +  50  ->  2HIO,. 


278  INORGANIC   CHEMISTRY 

In  both,  these  actions  the  initial  substances  (including  the  excess  of 
nitric  acid)  and  the  products,  with  the  exception  of  the  iodic  acid 
itself,  are  all  volatile.  When  the  solution  is  concentrated  by  evapora- 
tion, the  iodic  acid  crystallizes.  It  is  a  white  solid,  perfectly  stable  at 
ordinary  temperatures,  and  can  be  kept  indefinitely.  At  170°  it  begins 
to  give  off  water  vapor  (2HI03  <=±  H20  -f  I205),  leaving  the  pentoxide 
of  iodine.  The  latter  is  a  white  crystalline  powder  which  may  be 
raised  to  300°  before  it,  in  turn,  breaks  up,  giving  iodine  and  oxygen. 
In  aqueous  solution  iodic  acid  is  an  oxidizing  agent,  but  does  not 
part  with  its  oxygen  so  readily  as  do  chloric  acid  and  bromic  acid.  It 
oxidizes  hydrogen  iodide  in  dilute  solution : 

HI03  +  6HI  ->  3H20  +  3I2, 

all  the  iodine  being  liberated  In  this  respect  it  resembles  concen- 
trated sulphuric  acid  (p.  237).  Dilute  sulphuric  acid  shows  no 
oxidizing  qualities. 

Various  Acids  Derived  from  One  Anhydride. — Some  acids  are 
related  to  their  anhydrides  as  are  hypochlorous  acid  (p.  268)  and 
sulphurous  acid  (p.  71).  One  molecule  of  the  anhydride  combines 
with  one  molecule  of  water.  In  other  cases,  however,  the  proportion 
of  water  may  be  less  or  greater  than  this.  Thus  phosphoric  anhydride 
(P205)  takes  up  three  formula-weights  of  water  (p.  71).  Now  if 
periodic  acid  were  of  the  former  type  (H20, 12O7  =  2HI04),  its  formula 
would  be  HI04.  It  does  form  salts  of  this  type,  such  as  NaI04  and 
AgI04.  But  the  free  acid  is  a  deliquescent  solid  of  the  formula 
H5I06  (=  5H20, 12O7),  and  the  most  easily  prepared  salt  belongs  to  this 
type.  All  types  are  called  periodates,  however,  because  their  composi- 
tions are  all  founded  upon  the  same  anhydride.  The  latter  has  not 
itself  been  made.  We  usually  speak  of  various  acids  and  salts  as 
being  derived  from  the  same  anhydride,  the  word  "derived"  being  used 
in  a  figurative  and  not  a  literal  sense. 

The  difference  between  two  acids  HI04  and  H5I06  is  not  at  all  the 
same  as  between  HI03  and  HI04.  The  latter  would  represent  differ- 
ent stages  of  oxidation,  being  derived  from  I2O5  and  I207  respectively, 
and  accordingly  would  be  named  iodic  acid  and  periodic  acid.  The 
former  differ  only  by  2H20,  and  an  addition  or  subtraction  of  the  two 
elements  of  water  in  equivalent  quantities  is  neither  oxidation  nor 
reduction.  Hence  they  are  both  periodic  acids  (see  Phosphoric  acid). 


OXIDES   AND   OXYGEN   ACIDS   OF  THE   HALOGENS  279 

Periodates  and  Periodic  Acid.  —  Sodium  periodate  (NaIO4)  is 
found  in  Chili  saltpeter.  When  sodium  iodate  (NaI03)  is  dissolved 
along  with  sodium  hydroxide  in  water,  and  chlorine  is  passed  into  the 
mixture,  the  sodium  hypochlorite  formed  from  the  latter  oxidizes  the 
iodate  (NaI03  -f-  O  — »  NaI04).  But  the  somewhat  insoluble  salt  which 
crystallizes  out  is  Na2H8I06 : 

NaI03  +  O  +  NaOH  +  H20  ->  Na2H3I06. 

Other  salts  may  be  made  from  this  one. 

An  aqueous  solution  of  periodic  acid  is  obtained,  like  that  of 
chloric  acid  (p.  274),  by  the  action  of  sulphuric  acid  on  barium  period- 
ate.  A  white,  very  soluble  solid  (H5I06)  remains  when  the  liquid  is 
evaporated.  When  this  is  heated,  water  and  oxygen  are  both  given 
off,  and  iodine  pentoxide  (I2O5)  alone  remains. 

Chemical  Relations. —  The  compounds  of  the  halogens  with 
.metals  and  with  hydrogen  diminish  in  stability,  with  ascending  atomic 
weight  of  the  halogen,  in  the  order:  F(19),  01(35.5),  Br  (80),  I  (127). 
Each  halogen  will  displace  those  following  it  from  this  kind  of  combi- 
nation. In  the  case  of  the  oxygen  compounds,  the  order  of  stability 
is  just  the  reverse,  those  of  iodine,  for  example,  being  the  only  ones 
which  are  reasonably  stable.  The  order  of  displacement  in  such  com- 
pounds confirms  this  conclusion. 

Amongst  the  oxygen  acids  of  any  one  halogen,  those  containing 
most  oxygen  are  most  stable.  The  salts  are  in  all  cases  more  stable 
by  far  than  the  corresponding  acids. 

The  halogens  when  combined  with  metals  and  hydrogen  are  univa- 
lent  (HI,  KOI,  etc.).  It  is  clear,  however,  that,  when  united  with  oxy- 
gen, their  valence  is  higher.  The  maximum  is  shown  in  perchloric 
anhydride  (C1207),  where  chlorine  appears  to  be  heptavalent. 

The  formulae  of  the  acids  might  be  written  so  as  to  retain  the 
univalence : 

H-C1,     H-0-C1,     H-0-0-C1,     H-0-0-0-C1, 
H-O-O-O-0-C1. 

But  compounds  in  which  we  are  compelled  to  believe  that  two  oxygen 
units  are  united  are  usually  unstable  (see  Hydrogen  peroxide),  and  we 
should  expect  the  instability  would  be  greater  with  three  and  with  four 
units  of  oxygen  in  combination.  Here,  however,  the  reverse  state 
of  affairs  must  be  taken  account  of  in  our  formulae,  for  HC104  is  the 


280  INORGANIC   CHEMISTRY 

most  stable  of  the  chlorine  set.  This  reasoning,  together  with  the 
heptavalence  in  C1,O7,  leads  us  to  assume  the  valence  seven  in  per- 
chloric acid  (see  Periodic  system).  The  structural  formulae  (cf.  p.  224) 
of  some  of  these  substances  are  therefore  often  written  as  follows : 

0  O 

II  II 

H-C1,        H-O-C1,        H-O-C1=0,        Na-0-I  =  0. 

II  II 

0  O 

Exercises. —  1.   Assign  to  its  proper  class  (p.  187)  each  of  the 
actions  mentioned  in  this  chapter. 

2.  Knowing  that  potassium  fluosilicate  (K2SiF6)  is  insoluble,  how 
should  you  make  chloric  acid  (p.  265)  ? 

3.  Make  the  equation  for  the  interaction  of  chlorine  with  calcium 
hydroxide  in  hot  water  (p.  266).     How  should  you  make  zinc  chlorate 
from  zinc  hydroxide  (Zn(OH)a)  ? 

4.  How  should  you  make  pure  potassium  hypochlorite  from  hypo- 
chlorous  acid  (p.  268)? 

5.  On  what  circumstances  would  the  possibility  of  making  barium 
chlorate  by  action  of  chlorine  on  barium  hydroxide  depend  (p.  274)  ? 

6.  Make  the  equations  for :  (a)  the  preparation  of  potassium  bro- 
mate  ;  (b)  pure  aqueous  bromic  acid ;  (c)  the  interaction  of  iodine  with 
aqueous  potassium  hydroxide  in  the  cold,  and  when  heated. 

7.  Using  the  method  given  on  p.  274,  make  the  equations  for  the 
interactions  of  chlorine  dioxide  with  water,  and  with  aqueous  potas- 
sium hydroxide. 


CHAPTER   XVII 
DISSOCIATION  IN  SOLUTION 

THE  employment  of  interacting  substances  in  the  form  of  solutions 
is  so  constant  in  chemistry,  and  the  reasons  for  this  are  so  cogent,  that 
we  must  now  resume  the  discussion  of  the  subject  of  solution  (cf. 
p.  145). 

The  present  chapter  will  be  devoted  to  giving  the  proofs  that,  to 
speak  in  terms  of  the  molecular  hypothesis,  the  molecules  of  acids, 
bases,  and  salts  in  aqueous  solutions,  are  actually  dissociated  into  parts 
by  the  solvent.  This  will  be  shown  by  consideration,  successively,  of 
certain  peculiarities  in  the  chemical  behavior,  the  osmotic  pressures, 
the  freezing-points,  and  the  boiling-points  of  the  solutions  of  these 
substances.  We  shall  see  that  these  parts  coincide  in  composition 
with  the  radicals,  and  are  called  ions.  Finally,  the  principles  of 
chemical  equilibrium  will  be  applied  to  the  relations  of  the  ions  to 
that  proportion  of  the  molecules  which  has  remained  undissociated. 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution.  —  Acids  all  contain  hydrogen  (p.  93). 
In  aqueous  solution,  if  soluble,  they  are  sour  in  taste,  they  turn  blue 
litmus  red,  and  their  hydrogen  is  displaced  by  certain  metals  (p.  95), 
and  has  the  properties  of  a  radical.  By  the  last  statement  is  meant 
that  it  very  readily  exchanges  places  with  other  radicals  in  reversible 
double  decompositions  (p.  264).  Many  other  bodies,  like  sugar,  kero- 
sene, and  alcohol,  contain  hydrogen  also,  but  not  one  of  them  shows 
all  of  these  properties.  Again,  all  salts  are  made  up  of  two  radicals, 
and  the  reversible  double  decompositions  into  which  they  enter  with 
acids,  bases,  and  other  salts,  consist  in  exchanges  of  these  radicals. 
Other  substances  may  include  the  same  combination  of  atoms,  but  in 
their  actions  these  groupings  are  often  disregarded.  Thus,  sodium 
chloride  and  silver  nitrate  exchange  radicals  completely  (p.  13),  and, 
in  dilute  solution,  hydrogen  chloride  and  sodium  hydrogen  sulphate 
do  so  partially  (p.  180).  But  sodium  chloride  and  nitroglycerine 
C8H5(lSr03)8  do  not  interact  at  all.  The  latter  is  not  a  salt,  although  it 
contains  the  same  proportion  of  nitrogen  to  oxygen  as  does  any  nitrate. 

281 


282  INORGANIC   CHEMISTRY 

Furthermore,  it  is  chiefly  in  aqueous  solution  that  these  special 
properties  of  acids,  bases,  and  salts  become  apparent.  Their  behavior 
is  often  quite  different  in  the  absence  of  this  solvent.  If,  for  example, 
we  mix  ammonium  carbonate  and  partially  dehydrated  cupric  nitrate, 
and  apply  heat,  a  violent  interaction  begins.  An  immense  cloud  of 
smoke  and  gas  is  thrown  out  of  the  tube,  and  the  substance  remaining 
is  either  black  or  reddish,  in  parts,  according  to  the  proportions  of  the 
substances  employed.  The  residue  contains  cupric  oxide,  and  some- 
times red  cuprous  oxide  (Cu2O).  The  gas  is  tinged  red  by  the  pres- 
ence of  nitrogen  peroxide  (N02),  while  a  more  careful  examination 
would  show  that  it  contained  carbon  dioxide,  nitrogen,  nitrous  oxide 
(N20),  water  vapor,  and  perhaps  still  other  products.  The  contrast, 
when  the  substances  are  dissolved  in  water  before  being  brought  in 
contact  with  one  another,  is  very  great.  A  pale-green  precipitate  is 
formed  at  once,  and  rapidly  settles  out.  On  examination,  this  turns 
out  to  be  cupric  carbonate,  while  evaporation  of  the  solution  furnishes 
us  with  ammonium  nitrate.  There  are  only  two  main  products,  and 
the  essential  part  (a  basic  cupric  carbonate  is  really  formed)  of  the 
action  in  solution  is  represented  by  the  equation : 

(NH4)2C03  +  Cu(N03)2-*CuC03  +  2NH4N03. 

In  the  interaction  between  the  dry  substances  the  molecules  are  com- 
pletely disintegrated,  and  the  whole  change  is  very  complex.  In  the 
action  in  water  no  heating  is  required,  the  substances  are  neatly  broken 
apart,  certain  groups  of  atoms,  which  we  call  radicals,  are  transferred 
as  wholes  from  one  state  of  combination  to  another,  and  the  rearrange- 
ment takes  place  in  a  machine-like  manner.  Contrasts  like  this 
between  the  interactions  of  anhydrous  and  dissolved  bodies  are  very 
common.  Thus,  we  have  had  occasion  (p.  96)  to  mention  the  difference 
between  the  action  of  metals  on  concentrated  and  on  dilute  sulphuric 
acid. 

Many  compounds,  however,  do  not  show  any  change  in  behavior 
when  dissolved  in  water.  Sugar,  for  example,  is,  as  a  rule,  more  readily 
acted  upon  in  the  absence  of  any  solvent.  Then  again,  while  water 
is  not  the  only  solvent  which  has  the  effect  we  have  just  described, 
the  majority  of  solvents,  if  they  affect  chemical  change  at  all,  simply 
retard  it.  Thus  the  union  of  iodine  and  phosphorus  in  the  absence  of 
a  solvent  takes  place  spontaneously  with  a  violent  evolution  of  heat. 
When  the  elements  are  dissolved  in  carbon  bisulphide  before  being 
mixed  the  action  is  much  milder,  although  the  product  is  the  same 


DISSOCIATION  IN  SOLUTION  283 

(phosphorus  tri-iodide).  The  diminution  in  the  concentration  of  the 
ingredients  has  decreased  the  speed  of  the  action  in  the  normal  way 
(p.  249).  That  water  and  some  other  solvents  have  a  specific  influence 
tending  to  increase  the  activity  of  certain  classes  of  substances,  shows 
that  a  special  explanation  of  the  phenomenon  must  be  found. 

Summing  up  these  points  we  see  that  the  peculiarity  of  acids,  bases, 
and  salts  in  aqueous  solution  is  that  each  compound  always  splits  in 
the  same  way.  Thus,  cupric  nitrate  always  gives  changes  involving 
Cu  and  N08  and  never  interacts  so  as  to  use  CuN2  and  03,  or  Cu02  and 
N02,  as  the  basis  of  exchange.  Similarly,  acids  always  offer  hydrogen 
in  exchange,  and  so  nitric  acid  behaves  as  if  composed  of  H  and  N08, 
and  sulphuric  acid  as  if  composed  of  2H  and  S04,  and  never  as  if 
made  up  of  HSO  and  H03,  or  H2S  and  O4.  The  sour  taste  and  the 
effect  upon  litmus  seem  to  be  properties  of  this  easily  separable  hydro- 
gen, for  they  are  shown  only  by  acids.  The  result  is  that  we  can 
make  a  list  of  the  units  of  exchange,  such  as  H,  OH,  N08,  C08,  S04, 
Cu,  K,  and  Cl,  employed  by  acids,  bases,  and  salts  in  their  interac- 
tions. The  molecule  of  each  compound  of  these  classes  contains  at 
least  two  of  them.  Even  when  these  units  contain  more  than  one 
atom,  their  coherence  is  as  noticeable  within  this  class  of  actions,  as 
is  the  permanence  of  the  atomic  masses  themselves  in  all  actions. 

The  question  raised  in  our  minds  is  whether  solution  in  water 
alters  the  character  of  the  molecule  simply  by  producing  a  sort  of 
plane  of  cleavage  in  it  which  creates  a  predisposition  to  a  uniform 
kind  of  chemical  change,  or  whether  it  actually  divides  the  molecules 
into  separate  parts  consisting  of  the  above  units  of  exchange,  and 
leaves  subsequent  chemical  actions  to  occur  by  cross-combination  of 
these  fragments.  The  fact  that  the  dissolved  substances  can  be  recov- 
ered by  evaporation  of  the  liquid  does  not  demonstrate  that  they  have 
not  been  changed  temporarily  while  in  solution.  The  alteration  which 
the  water  produces,  whatever  it  be,  will  naturally  be  reversed  when  the 
water  is  removed.  Since  our  question  involves  nothing  but  the  count- 
ing of  particles,  the  number  of  which  would  be  much  greater  in  the 
event  that  actual  subdivision  of  molecules  is  the  explanation,  it  can 
be  answered  by  a  study  of  the  physical  properties  of  solutions. 

OSMOTIC  PRESSURE. 

In  the  earlier  discussion  of  solution  (p.  150)  the  condition  of  a 
dissolved  substance  was  viewed  as  akin  to  that  of  a  gas.  We  con- 
ceived the  molecules  of  the  dissolved  substance  as  being  distributed 


284 


INORGANIC    CHEMISTRY 


through  the  space  occupied  by  the  solvent,  as  being  separate  from  one 
another,  and  as  moving  about  independently  of  each  other.  This  was 
because  the  phenomena  of  diffusion  and  osmotic  pressure  (p.  151)  in 
solution  closely  resemble  those  of  diffusion  and  pressure  in  gases. 

The  invention  of  a  suitable  hypothesis  for  the  explanation  of  the 
facts  of  osmosis  presents  some  difficulties,  but  the  facts  themselves 
are  undoubted.  It  will  conduce,  therefore,  to  clear- 
ness if  we  speak  first  of  some  things  which  may 
be  observed  and  are  true,  irrespective  of  any 
explanation. 

Phenomena  Produced  by  Osmotic  Pres- 
sure. —  In  order  that  the  osmotic  pressure  (Gk. 
WO-/AOS,  impulsion)  of  the  molecules  of  a  diffusing 
body  may  be  perceived,  a  partition,  which  they 
are  unable  to  traverse,  must  be  interposed  between 
the  solution  and  a  contiguous  mass  of  the  pure 
solvent  (Fig.  57,  p.  151).  The  partition  must  be 
permeable  by  the  solvent,  however.  Such  a  parti- 
tion is  described  as  semi-permeableo 

The  general  nature  of  the  phenomena  may  be 
seen  by  employing  a  tube  (Fig.  73),  to  which  a 
diffusion-thimble  (Schleicher  and  Schull,  No.  579) 
is  securely  attached.  It  is  charged  with  sugar 
solution,  and  suspended  in  pure  water.  This  thim- 
ble is  somewhat  permeable  by  the  sugar,  but  the 
water  traverses  it  very  easily,  and  so  an  exhibition 
of  the  general  result  of  a  stricter  test  is  obtained 
/  \  quickly. 

Fio.  73.  The  water  is  able  to  pass  freely  through  the  mem- 

brane in  either  direction,  while  the  sugar  is  not.  As 
the  result  of  the  interchange  of  water,  the  liquid  rises  slowly  but 
steadily  in  the  tube.  The  pure  solvent  always  passes  into  the  solu- 
tion. If,  further,  two  solutions  of  different  concentrations  of  the  same 
substance  are  employed,  then,  invariably,  water  passes  from  the  more 
dilute  solution  into  the  more  concentrated  one  through  the  membrane. 
There  is  apparently  a  tendency  for  the  water  so  to  distribute  itself 
that  the  solutions  may  eventually  become  equal  in  strength.  The 
water  passes  from  a  dilute  solution,  leaving  it  more  concentrated  than 
before,  into  a  more  concentrated  solution,  rendering  it  more  dilute. 


DISSOCIATION  IN   SOLUTION  285 

These  phenomena  were  first  studied  by  Pfeffer  (1877),  a  botanist, 
who  used  certain  plant  cells  for  the  purpose.  The  cell  content  in- 
cluded a  liquid  containing  various  salts  in  solution,  and  a  protoplas- 
mic layer  which  was  not  attached  to  the  cell  wall.  This  protoplasmic 
layer  behaved  like  a  semi-permeable  membrane.  When  such  cells 
were  immersed  in  a  concentrated  solution  of  any  substance,  the  water 
passed  from  the  interior  of  the  cell  to  the.  solution,  and  by  means  of 
a  microscope  a  shrinkage  of  the  protoplasmic  layer  away  from  the  cell 
wall  could  be  observed.  Conversely,  when  such  cells  were  placed  in 
pure  water,  or  a  solution  of  a  very  dilute  nature,  water  passed  from  the 
outside  into  the  interior,  and  the  protoplasmic  layer  was  distended  so 
as  to  fill  the  corners  completely.  The  distension  of  the  cells  of  droop- 
ing flowers,  when  their  stems  are  placed  in  water,  and  the  consequent 
revival,  is  a  familiar  illustration  of  the  same  sort  of  thing.  All  solu- 
tions which  produced  neither  the  one  effect  nor  the  other  on  a  given 
set  of  plant  cells,  were  named  is-osmotic.  The  osmotic  pressures  of 
their  contents  were  the  same  as  the  pressure  of  the  cell  fluid. 

Professor  Crum  Brown  has  devised  an  arrangement  which  exhibits  the  action 
of  a  perfectly  seini-perineable  membrane  very  strikingly.  A  concentrated  solution 
of  calcium  nitrate  is  shaken  with  a  small  amount  of  phenol  (carbolic  acid),  so  as 
to  become  saturated  with  the  latter,  and  the  mixture  is  then  poured  into  a  tall, 
narrow  cylinder.  The  phenol  rises  and  floats  upon  the  surface  of  the  calcium 
nitrate.  The  amount  of  phenol  should  not  be  more  than  sufficient  to  saturate  the 
liquid  and  give  a  layer  a  few  millimeters  in  thickness.  Distilled  water,  also  satu- 
rated with  phenol,  is  cautiously  introduced  above  all.  The  water  on  both  sides  of 
the  layer  of  phenol  is  soluble  in  phenol,  and  consequently,  by  dissolving  in  this 
and  passing  out  on  the  other  side,  can  traverse  the  partition.  The  calcium  nitrate, 
however,  which  is  here  the  dissolved  substance,  cannot  traverse  the  phenol  in 
which  it  is  not  soluble.  The  phenol  therefore  constitutes  a  perfect  semi-permeable 
membrane.  If  the  level  of  the  lower  side  of  the  phenol  is  marked  on  the  outside 
of  the  cylinder  by  means  of  a  strip  of  paper,  it  will  be  found,  as  the  arrangement 
is  watched  from  day  to  day,  that  the  water  passes  through  the  phenol  into  the  solu- 
tion, and  the  phenol  rises  higher  and  higher,  until  finally  it  surmounts  all  the  rest 
of  the  liquid. 

The  Phenomena  a  Logical  Consequence  of  Semi-Permea- 
bility. —  The  passage  of  the  water  into  the  solution  .in  which  the 
greater  osmotic  pressure  exists  seems  at  first  paradoxical.  We  must 
remember,  however,  that  the  system,  consisting  of  the  liquids  on  each 
side  of  the  membrane,  can  be  in  equilibrium  only  when  the  osmotic 
pressure  on  the  two  sides  is  identical.  But  the  equalization  of  the 
osmotic  pressures  cannot  take  place  by  the  passage  of  part  of  the 
solute  from  one  side  to  the  other.  The  membrane  has  been  taken, 


286  INORGANIC   CHEMISTRY 

purposely,  of  such  a  nature  that  the  dissolved  substance  is  unable  to 
traverse  it.  The  equalization  must  occur,  therefore,  in  the  only  other 
possible  manner,  namely,  by  the  passage  of  the  solvent  in  the  other 
direction. 

An  imitation  of  this  behavior  may  easily  be  exhibited  by  the  use  of  gases.  A 
piece  of  peritoneal  membrane  is  stretched  across  the  mouth  of  a  thistle-tube  and 
moistened  with  water.  The  tube,  which  has  been  bent  in  U-f°rm  to  serve  as  a 
manometer,  contains  a  small  amount  of  some  colored  liquid,  whose  motions  will 
exhibit  any  change  in  pressure  in  the  interior.  When  an  inverted  cylinder  of 
ammonia  gas  is  placed  round  the  head  of  the  thistle-tube,  the  ammonia  gas  dis- 
solves in  the  water  on  the  membrane  until  this  water  is  saturated,  that  is,  until 
the  ammonia  molecules  leaving  the  water  are  as  numerous  as  those  entering 
it.  Itjwill  be  seen,  however,  that  the  ammonia  solution  really  has  two  surfaces, 
one  of  them  towards  the  interior,  and  the  ammonia  particles  must  eventually  leave 
both  surfaces  at  the  same  rate  at  which  they  are  landing  upon  one  of  them.  The 
ammonia  gas  being  at  the  pressure  of  the  atmosphere,  the  particles  of  ammonia 
leaving  the  film  will  produce  a  tension  of  one  atmosphere  of  ammonia  over  each 
surface.  Thus  ammonia  gas  will  be  transferred  from  the  cylinder  to  the  interior 
of  the  thistle-tube  until  its  partial  pressure  in  the  latter  is  equal  to  that  in  the 
former.  The  membrane  is  semi-permeable,  since,  of  the  air  and  ammonia  con- 
tained in  the  thistle-tube,  only  the  ammonia  can  traverse  the  film.  The  con- 
tents of  the  thistle-tube  therefore  correspond  to  the  solution,  air  being  the  solute 
and  ammonia  the  solvent.  The  original  air  in  the  apparatus  was  at  a  pressure  of 
one  atmosphere,  but  the  ammonia,  although  under  no  greater  pressure,  enters 
nevertheless.  Indeed,  it  would  continue  to  do  so  until  the  pressure  inside  became 
equal  to  that  of  the  ammonia  outside  plus  the  original  pressure  of  the  air,  a  total  of 
two  atmospheres.  The  case  corresponds  to  that  of  water  entering  a  solution  whose 
osmotic  pressure  is  one  atmosphere.  It  enters  until  the  contents  of  the  apparatus 
are  under  a  pressure  one  atmosphere  greater  than  that  existing  outside. 

Measurement  of  Osmotic  Pressure.  —  It  will  be  seen  that  the 
whole  phenomenon  rests  upon  the  fact  that  the  membrane  used  is 
permeable  by  one  of  the  constituents  only.  The  preparation  of  a 
vessel  of  sufficient  strength,  and  possessing  walls  with  the  maximum 
permeability  by  water  and  the  minimum  permeability  by  dissolved 
substances,  presents  great  difficulties.  A  device  of  Pfeffer's  is  still 
found  to  be  the  best.  A  cylinder  of  porous  porcelain,  much  like  a 
Pasteur  filter-tube,  is  treated  so  that  its  pores  are  partially  filled  with 
a  gelatinous  precipitate  of  cupric  ferrocyanide  (q.v.). 

The  porous  cylinder,  after  removal  under  the  air-pump  of  the  air  which  its  walls 
contain,  is  placed  in  a  solution  of  cupric  sulphate.  Its  interior  is  then  filled  with  a 
solution  of  potassium  ferrocyanide.  When  these  two  liquids  meet  by  diffusion 
inside  the  wall,  they  interact,  producing  a  dense  precipitate  of  the  substance  above 
mentioned  : 

2CuS04  +  K4Fe(CN)c  ->  Cu2Fe(CN)6  +  2K2SOV 


DISSOCIATION  IN  SOLUTION 


287 


If  such  a  prepared  vessel,  after  being  filled  with  a  one  per  cent 
sugar  solution,  could  be  closed  by  a  piston  (e.g.  Fig.  46)  and  be  placed 
in  pure  water,  it  would  be  found  necessary  to  place  weights  on  the 
piston  to  prevent  an  upward  movement,  due  to  access  of  water  to  the 
interior  through  the  walls.  Finally  a  weight  would  be  found  that  would 
just  balance  the  inward  tendency  of  the  water.  With  more  weight 
than  this,  water  would  be  squeezed  out  through  the  pores ;  with  less, 
the  water  would  force  its  way  in  and  the  piston  would  rise.  When 
this  weight  has  been  placed  in  position,  the 
water  inside  and  outside,  having  reached  a 
condition  of  equilibrium,  must  be  exerting 
equal  pressures  on  each  side  of  the  wall  of  the 
vessel.  Hence,  the  excess  of  pressure  inside 
must  be  due  to  the  osmotic  pressure  of  the 
dissolved  sugar.  It  cannot  be  due  to  the  water 
itself,  for  that  is  able  to  escape  through  the 
pores.  The  weight  opposing  the  osmotic  press- 
ure at  15°  in  the  case  of  a  one  per  cent  sugar 
solution  is  found  to  be  about  0.7  kg.  for  every 
sq.  cm.  of  the  exposed  surface.  Since  1.03  kg. 
per  sq.  cm.  equals  760  mm.,  this  would  indicate 
a  pressure  of  760  x  0.7  -H  1.03,  or  516  mm. 
(0.68  atmospheres). 

In  practice  a  small  bent  tube  opening  into 
the  cylinder  is  used  as  a  manometer  (Fig.  74). 
The  other  end  of  the  tube  is  closed,  and  some 
air  is  confined  in  this  end  by  mercury.  The 
diminution  in  the  volume  of  the  air  registers 
the  pressure.  The  smaller  tube,  drawn  out  to 
a  point,  is  used  for  filling  the  cell  with  the 
solution  and  is  then  sealed  before  the  blow- 
pipe. The  whole  apparatus  is  immersed  in 
a  large  bath  of  water  whose  temperature  can 

be  maintained  constant  during  the  experiment.  Concordant  readings 
are  hard  to  get  in  consequence  of  difficulties  inherent  in  the  prepara- 
tion and  use  of  the  apparatus,  but  the  general  relations  of  the  results 
can  be  stated  in  a  very  simple  form. 

Ten  years  after  Pfeffer's  experimental  work,  van  't  Hoff  first  formulated  the 
laws  of  osmotic  pressure.  He  showed  that  the  general  analogy  between  the  gase- 
ous state  and  the  state  of  solution  could  be  developed  so  as  to  exhibit  a  complete 


FIG.  74. 


288 


INORGANIC   CHEMISTRY 


correspondence  between  the  laws  of  both.  His  conclusions  (1887)  were  founded 
partly  on  Pfeffer's  results  and  partly  on  supplementary  experiments,  and  are  given 
in  the  three  following  paragraphs  (see,  also,  appendix  to  this  chapter). 

Osmotic  Pressure  and   Concentration.  —  A   part    of    one   of 
Pfeffer's  sets  of  experiments  will  show  the  relation  in  this  respect : 


PER  CENT  OF  SUGAR. 

OSMOTIC  PRESSURE. 

PER  CENT  OF  SUGAR. 

OSMOTIC  PRESSURE. 

1 
2 

535  mm. 
1016 

4 
6 

2082  mm. 
3075 

The  osmotic  pressures  of  a  series  of  solutions  of  the  same  substance 
are  proportional  to  their  concentrations.  The  form  of  the  law  is  the 
same  as  that  of  Boyle's  law  (p.  81).  The  pressures  are  in'general  the 
same  for  the  same  concentrations  whatever  solvent  is  used. 

Osmotic  Pressure  and  Temperature. — Pfeffer,  using  a  one  per 
cent  solution  of  sugar  at  two  different  temperatures,  found  the  pressure 
at  14.2°  to  be  510  mm.  and  at  32°  to  be  544  mm.  Not  only  does  osmotic 
pressure  change,  with  alteration  in  temperature,  in  the  same  direction 
as  does  gaseous  pressure,  but  the  changes  can  be  expressed  by  Charles' 
law  (p.  87).  The  osmotic  pressure  increases  in  proportion  to  the 
absolute  temperature.  A  gas  which  at  14.2°  C.  exhibits  a  pressure  of 
510  mm.,  at  32°  exercises  a  pressure  of  542  mm. 

An  Analogue  of  Avogadro  's  Hypothesis.  —  Still  more  interest- 
ing is  the  fact  that,  if  we  compare  the  concentrations  of  different  solu- 
tions which  at  the  same  temperature  exhibit  equal  osmotic  pressures, 
we  find  that  they  contain  equal  numbers  of  molecules  of  the  dissolved 
substance  in  equal  volumes.  Thus,  if  we  dissolve  one  mole  (342  g.) 
of  sugar  (C12H22On)  and  one  mole  (74  g.)  of  methyl  acetate  (CH3C2H802) 
in  equal  volumes  of  water,  we  have  taken  equal  numbers  of  molecules 
of  the  two  substances,  and  the  osmotic  pressures  which  the  solutions 
exhibit  are  found  to  be  equal. 

As  will  be  seen  below,  certain  substances  in  certain  solvents,  par- 
ticularly in  water,  exhibit  pressures  which  are  greater  than  this  hypoth- 
esis would  permit. 

The  closer  study  of  the  figures  enabled  van  't  Hoff  (1887)  to  state 
the  most  interesting  fact  of  all :  The  osmotic  pressure  exercised  by  a 
substance  in  solution  is  identical  in  value  with  the  gaseous  presssure 


DISSOCIATION  IN  SOLUTION  289 

which  it  would  exhibit  if  the  same  quantity  of  it  were  contained  as 
a  gas  in  the  same  volume  at  the  same  temperature.  For  example, 
44  g.  of  carbon  dioxide  in  the  gaseous  condition  fills  the  G.M.V. 
(22.41.),  and  at  0°  exercises  a  pressure  of  one  atmosphere.  We  find 
that  when  we  dissolve  the  same  quantity  of  the  same  substance  in 
22.4 1.  of  any  solvent  at  the  same  temperature,  it  exercises  one  atmos- 
phere of  osmotic  pressure.  Certain  substances,  however,  particularly 
when  dissolved  in  water,  exhibit  greater  pressures  than  this  (see 
below). 

Determination  of  Molecular  Weights.  —  It  is  evident  that  we  have 
here  an  experimental  method  which  may  be  used  in  measuring  the  molecular 
weight,  and  is  applicable  to  substances  which  cannot  be  converted  into  vapor.  All 
that  is  necessary  is  to  dissolve  a  weighed  amount  of  the  substance  in  a  known 
amount  of  water,  or  some  suitable  solvent,  and  by  means  of  the  apparatus 
described  above  to  measure  the  osmotic  pressure  at  some  fixed  temperature.  From 
the  result,  by  means  of  the  laws  corresponding  to  those  of  Boyle  and  Charles,  we 
may  calculate  the  concentration  of  the  solution  which  would  have  given  one  atmos- 
phere pressure  at  0°.  That  quantity  of  substance  which  would  give  this  concen- 
tration in  22.4  1.  of  the  solvent  is  then  the  molecular  weight  (c/.  p.  195).  The  time 
required  for  measurements  of  osmotic  pressure  and  the  experimental  difficulties 
alone  prevent  the  employment  of  this  method  in  practice. 

Osmotic  Pressure  and  Dissociation  in  Solutions.  —  What 
inference  is  to  be  drawn  in  the  cases  in  which  abnormally  high  osmotic 
pressures  are  observed  ?  In  view  of  the  fact  that  the  pressure  is  sup- 
posed to  be  produced  by  the  impact  of  the  particles,  and  depends  on 
the  number  of  them  in  the  given  volume,  we  must  infer  that  where  the 
pressure  is  greater,  more  particles  are  present  in  the  given  volume 
than  we  had  supposed.  In  other  words,  dissociation  of  the  original 
molecules  must  have  occurred.  This  phenomenon  is  observed  when- 
ever acids,  bases,  or  salts  in  aqueous  solution  are  under  observation. 
Thus  a  solution  of  sugar,  which  does  not  belong  to  these  classes,  con- 
taining 342  g.  in  the  G.M.V.,  exhibits  the  normal  osmotic  pressure  of 
one  atmosphere  at  0°.  A  solution  of  one  molecular  weight  of  potas- 
sium chloride  (74.5  g.)  in  the  same  volume  of  water,  however,  exhibits 
an  osmotic  pressure  of  about  1.88  atmospheres  at  0°.  The  greater 
pressure  must  be  due  to  the  fact  that,  although  the  number  of  mole- 
cules of  potassium  chloride  taken  is  the  same  as  in  the  case  of  sugar, 
the  number  of  actual  particles  whose  impacts  constitute  the  pressure  is 
greater,  —  is,  in  fact,  88  per  cent  greater.  Now  the  multiplication  of 
particles  from  potassium  chloride  molecules  can  occur  only  by  their 


290 


INORGANIC   CHEMISTRY 


dissociation  into  particles  of  K  and  Cl  by  a  chemical  change  rep- 
resented by  the  equation  KC1  ^±  K  +  Cl.  In  this  case,  seeing  that 
each  original  molecule  can  give  but  two  particles,  the  excess  of  press- 
ure indicates  that  0.88  (88  per  cent)  of  the  molecules  of  potassium 
chloride  have  been  broken  up.  Comparison  shows  that  the  degree  of 
dissociation  for  equi-molar  solutions  of  different  acids,  bases,  or  salts 
varies  widely.  For  the  same  substance,  it  is  always  relatively  greater 
in  dilute  than  in  concentrated  solutions. 

It  will  be  seen  that  we  have  thus  a  purely  physical  and  perfectly 
independent  confirmation  of  the  indications 
already  found  in  the  chemical  behavior  of 
substances  of  this  kind.  In  practice,  on 
account  of  the  experimental  difficulties, 
this  method  is  not  used  for  measuring  the 
degree  of  dissociation. 

DEPRESSION  IN  THE  FREEZING-POINT 
OF  A  SOLVENT. 

Measurement  of  Freezing -Point  8. 

—  The  task  consists  in  measuring  exactly 
the  temperature  at  which  a  previously 
weighed  quantity  of  the  solvent  freezes, 
and  then,  after  dissolving  in  it  a  known 
weight  of  some  soluble  substance,  deter- 
mining the  freezing-point  once  more.  The 
absolute  values  of  these  two  points  are  not 
required,  it  is  simply  the  difference  be- 
tween them  that  has  to  be  known  with 
exactness  (cf.  p.  163).  By  means  of  a 
very  delicate  thermometer  (Fig.  75)  having 
only  six  degrees  on  the  whole  scale,  the 
temperature  of  the  freezing  liquid  may 
be  read  to  one  one-thousandth  of  a  degree. 

A  reservoir  at  the  top  enables  us  to  add  to,  or  subtract  from,  the 
mercury  contained  in  the  bulb  and  column,  and  so  the  same  instrument 
may  be  used  with  solvents  having  widely  different  freezing-points. 
When  water  is  being  employed  as  the  solvent,  the  outer  jar  must  be 
filled  with  a  freezing  mixture  of  ice  and  water  containing  salt.  With 
solutions  in  benzene,  ice  and  water  are  used  alone.  To  avoid  super- 


FlG.  75. 


DISSOCIATION   IN   SOLUTION  291 

cooling,  the  solvent  or  solution  must   be  vigorously  stirred   after  it 
has  been  cooled  down  to  a  point  just  below  the  freezing-point. 

Laws  of  Freezing -Point  Depression.  —  The  depression  is 
directly  proportional  to  the  weight  of  dissolved  substance  in  a 
given  amount  of  the  solvent.  The  depression  is  inversely  propor- 
tional to  the  amount  of  solvent.  Thus,  if  we  double  the  concen- 
tration of  the  solution,  the  depression  in  the  freezing-point  is  doubled. 
Further,  equal  numbers  of  molecules  of  different  solutes  in  the  same 
quantity  of  solvent  give  equal  depressions.  Or,  in  other  words,  the 
depression  is  proportional  to  the  concentration  of  the  molecules  of 
the  solute.  Thus,  solutions  containing  342  g.  of  sugar  (C12H22OU),  or 
46  g.  of  alcohol  (C2H60),  or  74  g.  of  methyl  acetate  (CH8C2H3O2),  in  1000 
g.  of  water,  show  a  depression  below  the  freezing-point  of  water  of 
1.89°  in  each  case.  This  depression  produced  by  a  mole  of  the  solute 
in  1  1.  of  water  is  called  the  molecular  depression  constant  and  has  a 
different  value  for  each  solvent.  For  solutions  of  the  same  molecular 
concentration  in  benzene  the  depression  is  4.9°,  in  phenol  (carbolic 
acid)  7.5°.  Combining  these  facts  in  one  expression  : 

The  observed  depression  \  Wt.  of  Solute  1000 


in  an  aqueous  solution  )  Mol.  Wt.  of  Solute     Wt.  of  Solvent. 

For  other  solvents,  the  corresponding  value  of  the  depression  constant 
must  be  substituted  for  1.89°. 

These  principles  may  be  expressed  mathematically  in  a  form  which  is  con- 
venient for  use.  If  A  represent  the  depression  in  any  actual  experiment,  5  the  de- 
pression produced  by  one  molecular  weight  in  1000  grams  of  solvent,  W  the 
weight  of  the  substance,  M  its  molecular  weight,  and  g  the  weight  of  the  solvent  in 
grams,  then : 

Wx  1000 

A  =  5  x  — YF 

M  x  g 

In  the  case  of  water,  as  we  have  seen,  5  is  1.89°.    For  each  solvent  the  value  of  S 
must  be  determined  by  means  of  a  substance  of  known  molecular  weight. 

These  laws  describe  the  facts  most  exactly  when  the  solutions  are 
dilute.  They  hold  only  when  there  is  no  chemical  interaction  between 
solute  and  solvent.  Even  so,  however,  acids,  bases,  and  salts  dissolved 
in  water  present  many  apparent  exceptions  and  must  be  discussed 
separately. 


292  INORGANIC   CHEMISTRY 

Determination  of  Molecular  Weights.  —  When  the  depression 
constant  of  a  solvent  has  once  been  ascertained  by  means  of  a  sub- 
stance of  known  molecular  weight,  this  method  may  be  used  for  deter- 
mining the  molecular  weight  of  other  substances  which  are  soluble  in 
the  same  liquid.  All  the  other  factors  can  be  observed  and  substituted 
in  the  formula.  This  method  is  especially  useful  when  the  substance 
cannot  be  converted  into  vapor  without  undergoing  decomposition  (see 
Hydrogen  peroxide). 

Freezing -Points   and  Dissociation  in   Solution.  —  The    sub 

stances  which  present  the  most  conspicuous  exceptions  to  the  above 
rules  are  acids,  bases,  and  salts  in  aqueous  solution.  With  most  of 
these,  the  depression  produced  is  greater  than  we  should  expect  from 
the  concentration  of  the  solution.  Thus,  in  an  actual  experiment,  two 
equi-molar  solutions  were  compared.  One  contained  one  mole  (74  g.) 
of  methyl  acetate,  and  the  other  one  mole  (58.5  g.)  of  sodium  chloride, 
each  dissolved  in  2000  g.  (2  liters)  of  water.  The  freezing-points  ob- 
served, on  the  arbitrary  scale  of  the  thermometer,  were : 

Pure  water 3.580°        Pure  water 3.580° 

Solution  of  methyl  acetate  .     2.610°         Solution  of  salt    .     .     .     1.902° 

Depression 0.970°        Depression 1.678° 

0.970° 


Excess  depression  by  salt  0. 708° 

The  solution  of  methyl  acetate,  as  it  contained  only  0.5  moles  of  the 
solute  per  liter  of  water,  showed,  as  it  should  do,  about  half  the  aver- 
age molecular  depression  (1.89°,  p.  291).  This  is  typical  of  the  class  of 
substances  showing  normal  behavior.  Sugar,  alcohol,  and  hundreds 
of  other  substances,  in  solutions  of  the  same  molar  concentration, 
would  have  given  the  same  value. 

The  freezing-point  of  the  salt  solution,  however,  was  much  lower. 
If  this  solution  had  contained  the  same  concentration  of  dissolved 
particles  as  the  other  solution,  its  depression  would  have  been  0.970° 
likewise.  The  number  of  particles  must  therefore  have  been  greater 
than  we  should  have  expected  from  the  number  of  molecules  taken. 
In  other  words,  a  portion  of  the  molecules  of  the  salt  must  have  been 
broken  up,  and  the  excess  depression,  0.708°,  must  have  been  due  to 
the  extra  particles  produced  by  dissociation.  Now  sodium  chloride 
molecules  cannot  give  more  than  two  particles  each,  and  the  depression 
is  proportional  to  the  number  of  particles.  It  follows,  therefore,  that 
|^§,  or  0.732  (73.2  per  cent)  of  the  molecules  were  dissociated. 


DISSOCIATION   IN   SOLUTION  293 

This  result  is  typical  also.  Acids,  bases,  and  salts  of  which  one 
mole  is  dissolved  in  two  liters  of  water,  are  found  to  give  irregular 
values,  all  more  or  less  in  excess  of  0.970°.  Those  which  contain  but 
two  radicals,  like  sodium  chloride  (NaCl)  and  potassium  nitrate 
(KNOj),  give  values  between  0.970°  and  2  x  0.970°.  Substances  like 
calcium  chloride  (CaCl2)  and  sodium  sulphate  (Na2S04)  give  depres- 
sions approaching  three  times  the  normal  value:  their  molecules 
contain  three  radicals.  The  excess  depression  depends,  therefore, 
upon  the  number  of  particles  which  each  molecule  can  furnish,  and 
upon  the  proportion  of  all  the  molecules  which  is  dissociated  into 
these  fragments. 

In  the  case  of  an  acid,  base,  or  salt,  the  depression  is  not  strictly 
proportional  to  the  concentration.  Thus,  one  mole  of  salt  in  four  liters 
of  water  does  not  give  half  the  depression  of  the  two-liter  solution 
(0.839°)  but  somewhat  more  (about  0.844°).  The  same  method  of 
calculation  indicates,  therefore,  a  greater  degree  of  dissociation 
(about  79  per  cent)  in  the  more  dilute  solution  (see  Ionic  equilib- 
rium, below). 

Acids,  bases,  and  salts,  so  far  as  they  are  soluble  in  materials  like 
toluene,  benzene,  chloroform,  and  carbon  bisulphide,  exhibit  simply 
normal  depressions  in  these  solvents.  It  appears,  therefore,  that  dis- 
sociation does  not  take  place  in  many  solvents.  In  common  experi- 
ence it  is  encountered  only  in  solutions  in  water,  and,  perhaps, 
alcohol. 

Boiling -Point  s  and  Dissociation  in  Solution*  —  If  space  per- 
mitted, a  series  of  statements  might  be  made  in  regard  to  the  boil- 
ing-points of  solutions  (cf.  p.  162)  which  would  be  closely  parallel  to 
those  about  freezing-points.  The  boiling-point,  as  we  have  seen,  is 
elevated,  however,  by  the  introduction  of  a  foreign  body.  Thus,  when 
water  is  the  solvent,  one  mole  of  a  solute  in  1000  g.  of  the  solvent 
normally  raises  the  boiling-point  0.52°  (that  is,  from  100°  to  100.52°). 
But  acids,  bases,  and  salts  form  an  exception  to  this  rule,  as  before, 
and  the  excess  elevation  which  they  give  is  a  measure  of  the  degree  of 
dissociation. 

Comparison  of  the  Results  of  the  Three  Methods.  —  When  we 
measure  the  osmotic  pressure,  the  freezing-point  depression,  and  the 
elevation  in  the  boiling-point  of  the  same  solution,  and  calculate  the 
degree  of  dissociation  from  the  result  of  each  measurement,  we  find 


294  INORGANIC  CHEMISTRY 

that  the  values  obtained  are  usually  identical,  within  the  limits  of 
error  to  which  the  methods  are  liable.  Indeed,  the  theory  of  this 
subject,  developed  by  van't  Hoff,  enables  us  to  connect  the  osmotic 
pressure  by  a  mathematical  relation  with  the  other  two  phenomena,  and 
to  calculate  any  one  of  the  three  from  any  other. 

The  connection  between  the  three  sets  of  phenomena  cannot  be  explained 
here.  It  is  treated  in  all  works  on  Physical  Chemistry.  It  may  be  pointed  out, 
however,  that,  in  one  essential  respect,  experiments  in  osmotic  pressure,  and  in 
the  freezing  and  boiling  of  solutions,  are  all  alike.  The  perception  of  osmotic 
pressure  involves  a  partition  which  the  solvent  alone  can  pass,  and  the  osmotic 
pressure  for  a  given  solution  is  the  one  required  to  force  the  solvent  out.  In 
freezing  a  solution,  pure  ice  is  separated,  and  so  a  similar  extrusion  of  a  part  of  the 
pure  solvent  is  effected.  In  a  boiling  solution,  for  which  the  above  rules  hold,  the 
vapor  is  composed  of  the  pure  solvent,  and  the  solute  remains  behind.  The  rela- 
tion between  the  three  operations  lies  in  the  fact  that  in  each  case  the  same  thing, 
namely,  the  separation  of  a  part  of  the  solvent,  is  done.  Each  method  effects  this 
in  a  different  way.  But  the  expressions  representing  the  work  done,  in  terms  of 
the  factors  which  define  the  work  in  each  case,  can  be  equated  in  pairs  and  the 
required  relation  established.  Thus  the  molecular  depression  of  the  freezing-point, 
or  the  molecular  elevation  in  the  boiling-point,  as  we  have  defined  them,  is  equal 
to  0.002  T2  -T-  g,  where  T  is  the  absolute  temperature  of  the  freezing-  or  boiling- 
point,  and  q  is  the  heat  of  fusion  or  vaporization,  as  the  case  may  be.  Water,  for 
example,  freezes  at  273°  abs.,  and  its  heat  of  fusion  is  79  cal.  per  gram,  from  which 
the  calculated  molecular  depression,  0.002  x2732  -=-  79,  or  1.88°,  is  obtained.  Simi- 
larly, using  the  boiling-point,  373°  abs. ,  and  the  heat  of  vaporization,  537  cal.  per 
gram,  we  calculate  the  molecular  elevation  of  the  boiling-point  to  be  0.618°. 

It  ought  to  be  added  that  abnormally  small  osmotic  pressures,  freezing-point  de- 
pressions, and  boiling-point  elevations,  are  also  frequently  observed.  This  occurs, 
however,  almost  wholly  in  non-aqueous  solvents,  such  as  benzene.  It  is  shown 
particularly  by  substances  containing  oxygen,  and  is  even  noticed  in  the  case  of 
acids,  bases,  and  salts.  By  parity  of  reasoning  we  infer  that  in  these  cases  associa- 
tion (cf.  p.  242)  of  the  molecules  has  occurred,  and  that  the  physical  unit  of  the 
solute  in  these  solvents  is  larger  than  the  ordinary  molecule. 

THE  APPLICATION  OF  THESE  CONCLUSIONS  IN  CHEMISTRY. 

The  Constitution  of  Solutions  of  Acids,  Bases,  and  Salts.— The 

composition  of  solutions  which  are  normal  or  abnormal,  in  respect  to 
osmotic  pressure,  freezing-point,  and  boiling-point,  may  be  shown  thus  : 


SOLUTES. 

DISSOLVED  IN  WA- 

TEB,  ALCOHOL,  ETC. 

DISSOLVED  IN 
TOLUENE,  CHLOBO- 

FOBM,  ETC. 

Abnormal 

Normal 

Normal 

Normal 

DISSOCIATION  IN  SOLUTION  295 

It  appears  that  water  and  some  other  solvents  have  the  power  of 
breaking  up  the  molecules  of  acids,  bases,  and  salts  and  of  holding  the 
fragments  apart  from  one  another  and  hindering  their  reunion.  In 
consequence  of  this,  our  view  of  the  nature  of  an  aqueous  solution  of 
hydrogen  chloride  (HC1),  or  common  salt  (NaCl),  or  sodium  hydroxide 
(NaOH),  or  any  of  the  substances  of  the  classes  which  these  represent, 
may  now  be  stated  in  definite  terms.  Such  a  solution  contains,  besides 
undivided  molecules  of  the  solute,  at  least  two  other  kinds  of  material, 
H,  Na,*  Cl,  OH,  etc.,  which  result  from  the  breaking  up  of  the  molecules. 
We  shall  see  that  these  subdivisions  of  the  original  molecules  have 
distinct  physical  and  chemical  properties  of  their  own.  The  descrip- 
tions of  the  "  properties  "  of  the  solutions,  as  they  used  to  be  given  in 
chemistry,  were  really  a  confused  statement  of  the  properties  of  the 
different  components  of  a  mixture  of  molecules  and  their  fragments. 
Thus  the  indications  of  dissociation  found  in  the  chemical  behavior 
of  acids,  bases,  and  salts  (p.  282)  are  fully  confirmed  by  a  study  of 
the  physical  properties  of  their  solutions.f 

The  suggestion  that  the  multiplication  of  particles  takes  place  by  interaction 
of  the  salt  with  part  of  the  water,  NaCl  4-  H2O  <=±  NaOH  +  HC1,  resulting  in  the 
production  of  two  molecules  of  dissolved  matter  from  one.  is  open  to  several  fatal 
objections.  In  the  case  of  a  highly  dissociated  salt,  according  to  this  explanation, 
the  mixing  of  the  acid  and  base  in  dilute  solution  should  result  in  no  particular 
change  and  give  rise,  therefore,  to  no  development  of  heat.  But  the  heat  of  neutral- 
ization is  very  great  in  such  cases.  This  is  an  example  of  a  stochastic  hypothesis 
(p.  142),  be  it  noted,  and  its  verity  or  falsity  can  be  put  to  the  test  at  once.  Its 
inapplicability  is  further  seen  in  the  fact  that  it  cannot  explain  the  dissociation  of 
acids  and  bases  themselves. 

The  free  radicals,  of  whose  existence  we  have  thus  become  convinced, 
constitute  a  new  set  of  materials.  Thus  the  hydrogen  radical  of  acids, 
although  a  form  of  uncombined  hydrogen,  differs  totally  from  the  gas 
which  is  composed  of  the  same  material.  The  latter  has  no  sour  taste 
or  effect  upon  litmus.  It  is  very  slightly  soluble  in  water,  while  the 
hydrogen  radical  exists  as  a  separate  substance  only  in  solution.  Again, 
substances  with  the  composition  of  the  radicals  N03  and  S04  are  not 
known  at  all  except  in  solutions.  The  chief  peculiarity  of  these  sub- 
stances is  that  a  solution  cannot  be  made  which  contains  less  than  two 

*  The  objection  that  separate  atoms  of  sodium  could  not  remain  free  in  water, 
will  be  disposed  of  later. 

t  Recent  observations,  showing  that  in  some  cases  rapid  double  decompositions 
of  the  normal  kind  take  place  in  solutions  which  exhibit  no  physical  evidence  of  the 
existence  of  dissociation,  demonstrate  that  it  would  have  been  unsafe  to  infer  dis- 
sociation from  chemical  evidence  alone. 


296  INORGANIC   CHEMISTRY 

kinds  of  them  side  by  side.     A  niche,  therefore,  must  be  created  in  our 
molecular  hypothesis  to  receive  these  new  substances. 

Nomenclature :  The  Ionic  Hypothesis.  —  Our  chemical  mole- 
cules are  the  units  of  material  in  the  gaseous  condition  (p.  198).  Evi- 
dently smaller  units,  which  may  nevertheless  contain  more  than  one 
atom,  must  be  assumed  to  exist  in  solution.  These  units,  for  a  reason 
that  will  appear  later,  are  called  ions,  and  their  composition  corresponds 
to  that  of  the  radicals.  The  dissociation  of  molecules  into  ions  is 
named  ionization.  The  substances  of  the  three  classes  which  alone 
are  ionized  may  be  designated  ionogena.  Since  ions  are  discrete  parti- 
cles, they  are,  in  all  physical  respects,  molecules.  Thus  we  speak  of 
the  molecular  concentration  of  ionic  hydrogen,  just  as  we  do  of  that  of 
dissolved  or  of  gaseous  hydrogen. 

The  solution  of  an  ionized  substance  is  called  an  electrolyte  (</.•».),  and  often 
this  term  is  applied  also  to  acids,  bases,  and  salts  themselves,  because,  when  dis- 
solved, they  produce  electrolytes.  This  is  rather  a  confusing  metonymy,  however, 
because  these  bodies  by  themselves  are  not  conductors.  This  use  of  the  term  also 
introduces  obscurity  because  it  connects  the  ionization  with  electrolysis  and  always 
conveys  the  impression  that  the  latter  produces  the  former.  The  electrolytic  prop- 
erty of  ions  is  only  one  amongst  many  special  properties  of  electrolytes,  and  the 
majority  of  these  properties  are  chemical  and  have  nothing  to  do  with  electrolysis. 
Hence  we  have  preferred  the  more  general  word  "  ionogen." 

The  radicals  and  their  chemical  behavior  are  real,  and  all  the  pecul- 
iarities of  aqueous  solutions  of  acids,  bases,  and  salts  are  experimental 
facts.  Ions,  however,  like  corpuscles,  atoms,  and  molecules,  are  part 
of  our  great  system  of  formulative  hypotheses  and  are  added  to  it  in 
order  to  maintain  its  self -consistency.  We  apply  that  part  of  the 
hypothetical  system  known  as  Avogadro's  hypothesis  to  solutions,  and 
finding  portions  of  molecules  which  do  not  exist  in  the  gaseous  con- 
dition and  which  have  special  properties  of  their  own,  a  new  class  of 
unit  masses  has  to  be  established.  Molecules  are  units  which  are  not 
commonly  disintegrated  by  vaporization  (p.  198) ;  ions,  those  which 
are  not  commonly  disintegrated  in  double  decomposition  in  solution ; 
atoms,  those  which  are  not  commonly  disintegrated  in  any  chemical 
action.  But  there  are  exceptions  in  each  of  the  three  cases.  The  ionic 
hypothesis  was  first  suggested  by  Arrhenius  (1887)  immediately  after 
the  publication  of  van  ?t  Hoff's  correlation  of  the  facts  about  osmotic 
pressure  (p.  288). 

It  is  worth  noting  that  the  quantities  expressed  by  the  formulae  Al,  Ca,  and 
K,  when  existing  as  ions,  produce  equal  osmotic  pressures,  and  have  equal  effects 


DISSOCIATION   IN  SOLUTION  297 

upon  the  freezing-  and  boiling-points.  This  is  a  further  justification  for  our  choice 
of  chemical  unit  quantities  of  the  elements  (atomic  weights),  for  the  atomic  weights 
have  these  properties  in  common,  and  equivalents,  of  course,  do  not  (cf.  p.  210). 

Ionic  Equilibrium.  —  Since  the  ions  are  chemically  different  from 
their  parent  molecules,  their  formation  represents  a  variety  of  chemi- 
cal change.  The  change  does  not  involve  any  chemical  interaction  with 
the  water,  of  the  nature  of  hydrolysis,  for  cases  in  which  this  takes  place 
are  expressly  excluded  from  consideration.  It  is  simply  a  dissocia- 
tion, i.e.  reversible  decomposition  of  the  dissolved  substance. 

From  the  fact  that  the  proportion  of  molecules  ionized  is  shown  to 
become  greater  as  more  and  more  of  the  solvent  is  added,  and  that 
removal  of  the  solvent  diminishes  the  proportion  of  ions  to  molecules, 
and  finally  leaves  us  the  substance  entirely  restored  to  the  molecular 
condition,  we  know  that  this  is  a  reversible  action  and  therefore  a  true 
dissociation.  The  molecules  and  their  ions  adjust  themselves  like  the 
constituents  in  any  case  of  chemical  equilibrium.  In  the  cases  above 
mentioned  we  should  have  the  following  actions  taking  place  :  * 

HC1  <=>  H*  +  01'      NaCl  <±  Na  +  Cl'       NaOH  <=>  Na*  +  OH'. 

These  equilibria  are  all  of  precisely  the  same  nature  as  that  of 
phosphorus  pentachloride  vapor  (p.  255),  and  the  discussion  of  the 
latter  should  be  reexamined  and  applied  by  the  reader.  The  sole 
difference  is  that  here  change  in  volume  is  effected,  not  by  compression 
or  by  release  of  pressure,  but  by  removing  or  adding  water.  The 
adjustment  to  a  condition  of  equilibrium,  however,  seems  to  be  instan- 
taneous where  ions  are  concerned,  while  in  other  chemical  actions  it 
always  takes  a  perceptible,  and  often  a  considerable  interval  of  time. 

Using  Cv  Cy  and  Cz  for  the  molecular  concentrations  (numbers  of 
moles  per  liter)  of  the  molecules,  and  the  two  ions,  respectively,  we  have 
an  equilibrium  constant  (cf.  p.  254),  in  this  case  called  the  ionization 

constant : 

6  2  X  68 


When  we  dissolve  a  single  substance  which  gives  only  two  ions,  the 
molecular  concentrations  of  the  ions  are  necessarily  equal.     Hence, 

C2 
in  such  a  case,  -~  =  K.       When  some  other  ionogen  with  a  common 

ion  is  present,  however,  the  values  of  C2  and  C8  will  be  different. 

*  The  symbols  Na%  Cl',  etc.,  are  used  to  indicate  that  these  are  ions,  and  not 
identical  with  atoms,  Na,  Cl,  etc.  The  negative  radicals  and  hydroxyl  are  distin- 
guished thus,  NO3',  OH',  and  the  others  by  a  dot,  H',  K*. 


298  INORGANIC    CHEMISTRY 

Considering  the  form  of  the  above  mathematical  expression,  it  will 
be  seen  that  when  the  degree  of  ionization  is  great,  C2  and  (73  are  larger 
than  Cv  and  the  value  of  K,  the  ionization  constant,  will  be  great. 
On  the  other  hand,  in  the  case  of  feebly  dissociated  substances,  the 
value  of  K  will  be  small.  Furthermore,  if  by  the  addition  of  more 
water  we  diminish  all  the  concentrations,  this  will  momentarily  affect 
the  numerator  more  than  the  denominator  (cf.  p.  298).  In  order,  there- 
fore, that  the  value  of  the  whole  expression  may  remain  constant,  the 
concentrations  of  the  ions,  represented  by  (72  and  (78,  must  become 
greater,  and  can  only  do  so  at  the  expense  of  the  concentration  of  the 
undissociated  molecules,  represented  by  Cr  Our  formula,  therefore, 
represents  successfully  the  fact  that  dilution,  which  diminishes  the  con- 
centration of  all  the  substances,  produces  a  greater  degree  of  ionization. 

A  more  general  form  of  treatment  will  be  required  later.  If  a  be 
the  number  of  moles,  say  of  acetic  acid,  originally  taken,  v  the 
volume  of  the  solution  in  liters,  and  x  the  number  of  moles  ionized, 
then  the  molar  concentrations  at  equilibrium  will  be 

_,       a  —  x       ,    _  x 

C.  = and  (7,  =  Ca  =  - . 

v  v 

Therefore, 

/x\*     a  —  x  x2 

I-    -i =  K,    or, r-=J£ 

\vj  v  (a  —  x)v 

This  is  known  as  Ostwald's  dilution  formula. 

Exercises.  —  1.  A  one  per  cent  sugar  solution  gives  an  osmotic 
pressure  of  516  mm.  at  15°.  What  is  the  molecular  weight  of  sugar  ? 
Assume  that  the  sp.  gr.  of  the  solution  is  1. 

2.  What  gaseous  pressure  would  be  exerted  by  a  gas  of  the  same 
molecular  concentration  as  a  one  per  cent  solution  of  sugar  at  15° 
(p.  288)  ?     Compare  the  answer  with   the  osmotic  pressure  of  the 
solution. 

3.  What   depression   in  the   f.-p.  of  water  will  be  produced  by 
dissolving  10  g.  of  bromine  in  1  kg.  of  this  solvent  ? 

4.  What  depressions  in  the  f.-p.  of  benzene  and  of  phenol  would 
be  produced  by  10  g.  of  bromine  to  1  kg.  of  the  solvent,  if  no  chemi- 
cal action  took  place  ? 

5.  What   is  the  molecular   depression-constant    of    a  solvent   in 
which  5  g.  of  iodine  in  500  g.  of  the  solvent  lowers  the  f.-p.  0.7°? 

6.  What  is  the  degree  of  dissociation  of  zinc  sulphate  if  5  g.  of  it 


DISSOCIATION   IN   SOLUTION 


299 


dissolved  in  125  g.  of  water  produce  a  lowering  of  0.603°  in  the  f.-p.? 
What  is  the  molecular  concentration  of  each  of  the  three  substances 
present  in  this  solution  ? 

7.  What  will  be  the  approximate  b.-p.  of  a  solution  of  common 
salt,  saturated  at  100°  (p.  157)  ?  Assume  that  the  solute  is  80  per 
cent  dissociated. 


Appendix:  Recent  Measurements  of  Osmotic    Pressure. — 

Measurements  of  osmotic  pressure  by  Morse  and  Frazer,  published  since  the  fore- 
going was  written,  are  the  only  ones  yet  made  which  deal  with  pressures  of  more 
than  three  or  four  atmospheres. 

The  actual  measurement  is  made  according  to  the  general  method  described  on 
p.  287,  that  is,  by  means  of  a  porous  porcelain  cup  or  cylinder  which  contains 
the  semi-permeable  membrane  within  its  walls,  and  is  attached  to  a  manometer 
containing  an  inclosed  volume  of  air  over  mercury. 

The  membrane,  however,  is  deposited  electrolytically,  instead  of  by  diffusion. 
The  method  of  procedure  is  briefly  as  follows  :  The  porous  cylinder,  after  the 
removal  of  the  air  from  its  walls  by  "electrical  endosmose,"  is  surrounded  by  a 
copper  electrode,  and  both  cup  and  electrode  are  immersed  in  a  0. 1 N  solution  of 
copper  sulphate.  The  other  electrode,  the  anode,  is  placed  within  the  cup,  which 
is  filled  with  a  0.1  N  solution  of  potassium  ferrocyanide.  A  current  of  electricity 
with  an  electromotive  force  of  110  volts  is  then  passed  through  the  solutions  and 
the  porous  wall  of  the  cylinder,  from  the  copper  to  the  platinum  electrode,  and  the 
membrane  of  copper  ferrocyanide  is  deposited  either  upon  the  interior  surface  of 
the  cup,  or  within  its  walls.  By  proceeding  in  this  way,  only  from  one  to  three 
hours,  instead  of  several  days,  are  required  for  the  formation  of  a  suitable  membrane. 

The  following  table  contains  the  results  of  some  measurements  made  with 
solutions  of  cane-sugar  in  water.  The  proportionality  between  concentration  and 
osmotic  pressure  is  evident. 


MOLES  SUGAR  IN 
1000  GRAMS  WATER. 

OSMOTIC  PRESSURE 
(ATMOSPHERES). 

MOLES  SUGAR  IN 
1000  GRAMS  WATER. 

OSMOTIC  PRESSURE 
(ATMOSPHERES). 

0.2 
0.4 

4.83 
9.72 

0.5 
1.0 

12.16 
24.46 

From  data  already  obtained  by  Morse  and  Frazer,  it  seems  very  probable  that 
the  statement  of  one  of  van't  Hoff's  laws  (p.  288)  will  have  to  be  modified,  for  it  is 
found  that  "cane-sugar,  dissolved  in  water,  exerts  an  osmotic  pressure  equal  to 
that  which  it  would  exert  if  it  were  gasified  at  the  same  temperature  and  the  volume 
of  the  gas  were  reduced  to  that  of  the  solvent  in  the  pure  state"  (Amer.  Chem. 
Jour.,  July,  1905). 


CHAPTER  XVIII 
OZONE   AND   HYDROGEN    PEROXIDE 

A  FRESH,  penetrating  odor,  resembling  that  of  very  dilute  chlorine, 
was  noticed  by  van  Marum  (1785)  as  being  perceptible  near  an  elec- 
trical machine  in  operation.  Schonbein  (1840)  showed  that  the  odor 
was  that  of  a  distinct  substance,  which  he  named  ozone  (Gk.  o£«v, 
to  smell),  and  he  discovered  a  number  of  ways  of  obtaining  it.  It  is 
very  questionable  whether  there  is  any  ozone  in  the  air,  excepting 
temporarily  in  the  immediate  neighborhood  of  a  natural  or  artificial 
discharge  of  electricity. 

Preparation  of  Ozone.  —  The  most  satisfactory  way  of  prepar- 
ing ozone  (Og)  is  to  allow  electric  waves  to  pass  through  oxygen. 
The  apparatus  (Fig.  76)  consists  of  two  co-axial  glass  tubes,  between 
which  the  oxygen  flows.  The  waves  are  generated  by  connecting  an 


FIG.  76. 

outer  layer  of  tinfoil  on  the  outer  tube,  and  an  inner  layer  of  tinfoil 
in  the  inner  tube  with  the  poles  of  an  induction  coil.  With  dry,  cold 
oxygen,  about  7.5  per  cent  of  the  gas  is  easily  turned  into  ozone. 
Under  the  best  conditions  this  proportion  cannot  be  much  exceeded. 

Ozone  is  found  in  the  oxygen  generated  by  electrolysis  of  dilute 
sulphuric  acid  (p.  95).  Some  of  it  is  produced  when  sulphuric  acid 
acts  upon  oxides  which,  with  this  reagent,  liberate  oxygen,  e.g., 
2Ba02  -f  2H2S04  — >  2BaS04  +  2H20  +  02.  It  arises  during  the  slow 
oxidation  of  phosphorus  by  the  air,  resulting,  probably,  from  the  de- 
composition of  unstable,  highly  oxidized  bodies  which  are  formed 
during  the  action.  Oxygen  containing  as  much  as  15  per  cent  of  it 
is  produced  by  the  interaction  of  fluorine  and  water  (p.  241). 

300 


OZONE  AND  HYDROGEN  PEROXIDE  301 

Physical  Properties  of  Ozone.  —  Ozone  is  a  gas  of  blue  color. 
It  boils  at  —  119°,  so  that  when  a  mixture  of  oxygen  and  ozone  is  led 
through  a  U-tube  immersed  in  liquid  oxygen  (—  182.5°),  the  ozone  is 
liquefied.  The  opaque,  deep-blue  fluid  contains  only  about  14  per 
cent  of  oxygen,  and  this  may  be  removed  by  evaporation. 

Ozone  is  much  more  soluble  in  water  than  is  oxygen.  Its  solubility 
shows  that,  at  12°,  100  volumes  of  water  would  dissolve  50  volumes  of 
the  gas  at  one  atmosphere  pressure.  Its  solubility,  when  mixed  with 
oxygen,  is  in  proportion  to  its  partial  pressure  (p.  155). 

Chemical  Properties  of  Ozone.  —  Ozone  is  relatively  stable  only 
when  mixed  with  much  oxygen.  Hence  its  density  and  molar  weight 
cannot  be  ascertained  save  by  indirect  means.  The  weight  of  a  liter 
of  the  mixture  at  0°  and  760  mm.  having  been  measured,  the  ozone 
may  be  removed  by  absorption  in  turpentine  and  the  proportion  of  it 
present  in  the  gaseous  mixture  be  thus  ascertained.  For  example,  if  the 
weight  of  1  1.  was  1.468  g.  and  50  c.c.  were  absorbed  by  turpentine, 
there  were  950  c.c.  of  oxygen.  The  weight  of  this  oxygen  is  1000  : 
950  :  :  1.429  :  x,  from  which  x  =  1.361  g.  The  rest  of  the  weight, 
1.468-1.361  or  0.107  g.,  was  that  of  50  c.c.  of  ozone.  The  weight  of 
1  1.  of  ozone  at  0°  and  760  mm.  is  therefore  2.140  g.  The  molecular 
weight  (weight  of  22.4  1.)  is  thus  47.9  g.,  or  nearly  48  g.  The  for- 
mula of  ozone  is  therefore  03. 

When  ozonized  oxygen  is  heated,  the  ozone  is  decomposed  at  about 
250-300°.  The  action  for  its  formation  : 


is  therefore  reversible.  That  this  equation,  showing  that  three  mole- 
cules of  oxygen  give  two  molecules  of  ozone,  is  correct,  may  be  demon- 
strated by  measuring  the  diminution  in  volume  which  accompanies  the 
action.  If  a  shrinkage  of  5  c.c.  is  observed  in  forming  the  ozone,  it  is 
found  that  10  c.c.  more  are  then  absorbed  by  turpentine.  Thus  the  ozone 
occupied  10  c.c.,  and  the  total  oxygen  from  which  it  was  made  was 
therefore  15  c.c.  Hence  three  volumes  of  oxygen  give  two  of  ozone. 

The  formation  of  ozone  absorbs  much  energy  from  the  electric  waves, 
or,  in  other  methods  of  making  it,  from  the  concomitant  chemical 
changes  : 

0  +  02  =  08  -  32,400  cal. 

Ozone  is  a  much  more  active  oxidizing  agent  than  oxygen.  Mercury 
and  silver,  which  are  not  affected  by  the  latter,  are  converted  into 


302  INORGANIC   CHEMISTRY 

oxides  by  the  former.  Silver  gives  the  peroxide,  Ag202.  Paper  dipped 
in  starch  emulsion  containing  a  little  potassium  iodide  is  used  as  a 
test  for  ozone  : 

03  +  2KI  +  H20  ->  02  +  2KOH  +  I3. 

The  iodine  gives  a  deep-blue  color  to  the  starch  (cf.  p.  235).  This 
test,  however,  will  not  distinguish  ozone  from  chlorine  or  hydrogen 
peroxide,  and  may,  therefore,  be  used  only  in  the  absence  of  these  sub- 
stances. The  last  substance  is  always  present  in  the  air,  and,  since  air 
usually  shows  the  above  action,  is  probably  responsible  for  the  belief 
that  air  contains  ozone.  The  action  on  silver  has  never  been  obtained 
with  air.  Ozone  also  removes  the  color  from  organic  dyes,  such  as 
indigo,  by  oxidizing  them  (cf.  p.  269).  Its  activity  as  an  oxidizing 
agent,  like  the  similar  activity  of  hypochlorous  acid,  is  due  to  the  fact 
that  it  contains  much  more  energy  than  oxygen.  In  all  its  actions  the 
energy  set  free  is  greater  by  this  excess  than  that  liberated  when  oxy- 
gen is  used. 

Oxygen  and  ozone  are  different  substances  (p.  35),  that  is,  have 
different  properties.  The  difference  in  density,  interpreted  in  terms 
of  the  molecular  hypothesis,  gives  us  the  statement  of  the  nature  of 
the  difference  which  is  embodied  in  the  formulae  02  and  08.  The  differ- 
ence in  activity,  interpreted  in  terms  of  the  conception  of  energy,  gives 
us  the  other  method  of  stating  the  nature  of  the  difference.  The  re- 
cent preference  for  the  second  method  is  well  illustrated  by  this  case. 
The  first  method  uses  a  mere  physical  property,  the  second  a  fact 
which  is  intimately  connected  with  the  whole  chemical  behavior  of  the 
substance,  a  matter  of  much  greater  interest  to  the  chemist. 

Ozone  may  be  distinguished  from  chlorine,  nitrogen  peroxide,  and  other  oxi- 
dizing agents,  with  the  exception  of  hydrogen  peroxide,  by  using  pink  litmus  paper 
instead  of  plain  paper  to  carry  the  potassium  iodide  solution  in  the  above  test. 
The  potassium  hydroxide  set  free  by  ozone  turns  the  paper  blue.  Chlorine,  for 
example,  gives  an  entirely  different  action  :  C12  +  2KI  — ->  2KC1  +  la- 
Ozone  is  used  commercially  in  bleaching  oils  and  in  purifying 
starch.  It  is  employed  also  for  sterilizing  drinking  water  in  Lille  and 
other  cities. 

HYDROGEN  PEROXIDE. 

Hydrogen  peroxide  (H202)  is  found  in  minute  amounts  in  rain  and 
snow.  It  is  formed  in  small  quantities,  in  a  way  not  at  present  under- 
stood, when  moist  metals  rust. 


OZONE  AND  HYDROGEN  PEROXIDE  303 

Preparation  of  Hydrogen  Peroxide.  —  When  sodium  peroxide 
(y.v.)  is  added,  a  little  at  a  time,  to  a  dilute  acid,  hydrogen  peroxide  is 
set  free : 

Na202  +  2HC1  t;  2NaCl  +  H202. 

It  may  be  separated  from  the  salt  (and  a  large  part  of  the  water) 
by  repeatedly  shaking  the  mixture  with  ether  (cf.  p.  155).  The  rela- 
tive solubility  in  water  and  ether  is  1 :  0.0596,  however,  so  that  much 
ether  is  needed.  The  ethereal  layer,  which  rises  to  the  top,  when 
evaporated,  leaves  a  strong  aqueous  solution  of  the  compound  behind. 

When  hydrated  barium  peroxide  (Ba02,  8H20)  is  shaken  with  cold, 
dilute  sulphuric  acid  a  similar  action  takes  place : 

Ba02  +  H2S04  £=»  BaS04  J  +  H202. 

The  excess  of  sulphuric  acid  may  be  removed  by  adding  barium 
hydroxide  solution  cautiously  until  no  further  precipitation  of  barium 
sulphate  occurs  :  Ba(OH)2  -f  H2S04  s=?  BaSOJ  +  2H20.  Hydrochloric 
acid  or  phosphoric  acid  may  be  used  instead  of  sulphuric  acid.  The 
second  is  largely  employed  in  the  commercial  manufacture  of  hydrogen 
peroxide.  In  each  case,  great  care  has  to  be  taken  to  precipitate  the 
other  products  and  all  impurities  from  the  solution.  When  hydro- 
chloric acid  is  used,  for  example,  the  barium  chloride  produced  by  the 
action  is  removed  by  adding  silver  sulphate  : 

BaCl2  +  Ag2S04  *=»  BaS04  J  +  2AgCl  J. 

An  aqueous  solution  is  also  obtained  by  passing  carbon  dioxide 
through  barium  peroxide  suspended  in  water : 

Ba02  +  C02  +  H20  ±=>  BaC03  J  +  H202. 

Pure  hydrogen  peroxide  is  isolated  from  any  of  these  solutions  by 
distillation  under  reduced-pressure  (p.  276).  It  is  much  less  volatile  than 
water,  but  decomposes  into  water  and  oxygen  violently  at  100°.  Hence 
the  lower  pressure  is  required  to  make  possible  its  volatilization  at  a 
temperature  below  this  point.  At  68  mm.  pressure,  the  water  begins  to 
pass  off  first  (at  about  45°).  The  last  portion  of  the  liquid  boils  at  84- 
85°  and  is  almost  all  hydrogen  peroxide. 

By  evaporating  the  commercial  (3  per  cent)  solution  at  70°,  a  liquid 
containing  45  per  cent  of  hydrogen  peroxide  may  be  made  without 
much  loss  of  the  material  by  volatilization. 


304  INORGANIC  CHEMISTRY 

The  Interaction  of  Barium  Peroxide  and  Sulphuric  Acid.  — 

It  is  worth  noting  that,  although  common  barium  peroxide  is  not  less 
soluble  in  water  than  is  the  hydrated  form,  it  dissolves  much  more 
slowly.  The  fact  that  it  is  made  by  heating  barium  oxide  in  oxygen 
and  is  composed  of  compact  particles  is  accountable  for  this. 

Every  action  upon  a  little-soluble,  or  slowly  dissolving  body,  like 
the  barium  peroxide  in  the  above  actions,  is  rather  complex.  It  is 
only  the  dissolved  part  of  the  substance  that  interacts.  There  is  thus 
a  physical  equilibrium  between  the  undissolved  and  the  dissolved 
bodies,  Ba02  (solid)  <=±  Ba02  (diss'd),  the  displacement  of  which  fur- 
nishes the  material  for  the  chemical  action.  The  latter  has  therefore  to 
follow  the  pace  set  by  the  former.  When  barium  sulphate  is  precipi- 
tated, another  physical  equilibrium  follows  the  chemical  change : 
BaSO4  (diss'd)  ^±  BaS04  (solid).  When  relatively  insoluble  bodies  are 
used  or  produced,  there  is  thus  a  chain  of  equilibria  each  depending 
on  the  others  : 

BaO2  (solid)  ±3  BaO2  (diss'd)  +  H2SO4 1^  H2O2  +  BaSO4  (diss'd)  ±^  BaSO4  (solid). 

If  the  barium  sulphate  ceased  to  be  precipitated,  its  interaction  in 
solution  with  the  hydrogen  peroxide  would  drive  the  central  action 
backwards,  and  barium  peroxide  would  be  precipitated  instead.  The 
success  of  the  process  thus  depends  on  the  fact  that  barium  sulphate 
is  even  less  soluble  than  barium  peroxide. 

When  carbon  dioxide  is  used  (see  above),  a  similar  chain  of  equili- 
bria exists,  and  in  that  case  it  is  the  barium  carbonate  that  is  the  less 
soluble  substance. 


Other  Modes  of  Formation.  —  Hydrogen  peroxide  is  formed  by 
the  direct  union  of  hydrogen  and  oxygen.  When  a  hydrogen  flame  is 
allowed  to  play  upon  ice,  appreciable  amounts  of  the  peroxide  are 
saved  from  being  decomposed,  as  they  ordinarily  would  be  by  the 
heat  of  the  action,  and  are  found  in  the  water. 

It  may  be  obtained  by  the  action  of  acids  upon  the  peroxides  of 
calcium,  strontium,  zinc,  and  copper. 

Traces  of  hydrogen  peroxide  are  formed  when  zinc,  copper,  lead,  and  other 
metals  are  shaken  with  air  and  dilute  sulphuric  acid.  It  is  produced  when  oxygen 
is  passed,  in  the  neighborhood  of  the  negative  electrode,  through  the  liquid  in  an 
electrolytic  cell  containing  dilute  sulphuric  acid.  The  gas  is  reduced  by  the  hydro- 
gen being  liberated  on  the  platinum  plate. 


OZONE  AND  HYDROGEN  PEROXIDE  305 

Physical  Properties.  —  Hydrogen  peroxide  is  a  syrupy  liquid  of 
sp.  gr.  1.5.  It  blisters  the  skin,  and,  when  diluted,  has  a  disagreeable 
metallic  taste.  It  has  been  frozen  (m.-p.  —  2°). 

Chemical  Properties.  —  Hydrogen  peroxide  is  very  unstable,  and 
decomposes  slowly  even  at  —  20°.  The  dilute  aqueous  solution,  when 
free  from  impurities,  keeps  fairly  well.  The  presence  of  a  trace  of 
free  acid  increases  its  stability.  Free  alkalies  and  most  salts  assist 
the  decomposition  ;  hence  the  necessity  for  purifying  the  commercial 
solution.  Addition  of  powdered  metals,  of  manganese  dioxide,  and 
of  charcoal  causes  effervescence  even  in  dilute  solutions,  and  oxygen 
escapes : 

2H202  -»  2H20  +  02. 

The  more  concentrated  solutions  (38  per  cent)  remain  quiescent  in  a 
dish  of  polished  platinum  even  at  60°,  but  the  making  of  a  slight 
scratch  on  the  bottom,  beneath  the  surface  of  the  liquid,  causes  pro- 
fuse liberation  of  oxygen  along  the  sharp  edge  thus  produced.  The 
action  of  the  catalytic  agents  is  therefore  probably  mechanical. 

Since  the  substance  cannot  be  vaporized,  even  at  low  pressure, 
without  some  decomposition,  its  molar  weight  has  been  determined  by 
the  freezing-point  method  (p.  291).  The  freezing-point  of  a  3.3  per 
cent  solution  in  water  was  2.03°  below  that  of  the  water  itself.  Hence, 
in  1000  g.  of  water,  3.3  g.  would  have  given  a  depression  of 
2.03  x  96.7  -f- 1000,  or  0.196°.  Therefore  a  depression  of  1.89°  would 
have  been  caused  by  3.3  x  1.89  -r-  0.196,  or  31.8  g.,  which  is  the  re- 
quired molar  weight.  Now  the  formula  HO  corresponds  to  a  molar 
weight  of  17  and  H202  to  one  of  34.  It  is  evident,  therefore,  that  the 
latter  is  the  correct  formula. 

Hydrogen  peroxide,  in  solution  in  water,  is  a  feeble  acid.  The 
normal  molar  weight  and  very  small  electrical  conductivity  (see  Chap, 
xix)  show  that  only  a  very  small  proportion  of  it  can  be  ionized.  As  an 
acid  it  enters  into  double  decomposition  readily.  Thus,  when  it  is 
added  to  solutions  of  barium  and  strontium  hydroxides,  the  hydrated 
peroxides  appear  as  crystalline  precipitates  : 

Sr(OH)2  +  H202  <=>  2H20  +  Sr02. 

The  precipitation  involves  another  equilibrium  :  Sr02  +  8H20  ^±  Sr02, 
8H2O  (solid).  The  action  of  hydrogen  peroxide  upon  chromic  acid 
(H2Cr04)  is  probably  of  a  similar  nature.  The  composition  of  the 
product,  which  gives  a  beautiful  blue  solution,  is  not  definitely  known, 


306  INORGANIC   CHEMISTRY 

as  it  decomposes  almost  immediately.  Its  formation,  by  adding  a 
drop  of  potassium  dichromate  to  an  acidulated  solution  of  the  peroxide, 
is  used  as  a  test  for  the  latter.  The  acid  interacts  with  the  dichro- 
mate, giving  the  necessary  chromic  acid  : 

H,S04  +  K2Cr207  -|-  HjjO  <=±  2H2Cr04  +  K2SO4. 

The  blue  substance  has  the  property,  unusual  in  inorganic  compounds, 
of  dissolving  much  more  readily  in  ether  than  in  water.  It  is  also 
much  less  unstable  when  removed  from  the  foreign  materials  in  the 
aqueous  solution.  Hence  the  test  is  rendered  more  delicate  by  ex- 
tracting the  solution  with  a  small  amount  of  ether.  In  the  ethereal 
layer  the  color  of  the  compound  is  more  permanent,  as  well  as  mor,e 
distinctly  visible  on  account  of  the  greater  concentration. 

Hydrogen  peroxide  is  a  much  more  active  oxidizing  agent  than  free 
oxygen.  It  liberates  iodine  from  hydrogen  iodide,  an  action  which, 
in  presence  of  starch  emulsion  (cf.  p.  235),  is  used  as  a  test  for  its 
presence  : 


It  converts  sulphides  into  sulphates.  The  white  lead  (q.v.)  used  in 
paintings  is  changed  by  the  hydrogen  sulphide  in  the  air  of  cities  to 
black  lead  sulphide,  PbC03  +  H2S  ->PbS  +  H20  +  CO2.  This  may 
be  oxidized  to  white  lead  sulphate  by  means  of  hydrogen  peroxide  : 

PbS  +  4H202  -4  PbS04  +  4H20, 

and  in  this  way  the  original  tints  of  the  picture  may  be  practically  re- 
stored. Organic  coloring  matters  are  changed  into  colorless  substances 
by  an  action  similar  to  that  of  hypochlorous  acid  (cf.  p.  269).  Hence 
hydrogen  peroxide  is  used  for  bleaching  silk,  feathers,  hair,  and  ivory, 
which  would  be  destroyed  by  the  more  violent  agent.  The  products 
of  its  decomposition,  being  water  and  oxygen  only,  are  harmless,  and, 
on  this  account,  it  is  used  as  a  bactericide  in  surgery. 

Hydrogen  peroxide  exercises  the  functions  of  a  reducing  agent  in 
special  cases,  also.      Thus,  silver  oxide  is  reduced  by  it  to  silver  : 

Ag20  +  H202  ->  2Ag  +  H20  +  Os. 

A  solution  of  potassium  permanganate,  in  which  the  permanganic  acid 
has  been  set  free  by  an  acid,  KMnO4  +  H2S04  +±  HMn04  +  KHSO4, 
is  rapidly  reduced.  The  permanganic  acid,  with  excess  of  sulphuric 
acid,  tends  to  undergo  the  first  of  the  following  changes,  provided  a 


OZONE  AND  HYDROGEN  PEROXIDE  307 

substance   is  present   which   can  take  possession  of  the   oxygon  that 
would  remain  as  a  balance : 

2HMn04  +  2H2S04  ->  2MnS04  +  3H20  (+  50)    (1) 

(50)  +  5H202  -» 5H20  +  502 (2) 

2HMn04  +  2H2S04  +  5H.O,  ->  2MnS04  -f  811,0  +  50, 

The  first  partial  equation  has  been  doubled  to  secure  the  even  num- 
ber of  units  of  hydrogen  required  for  the  formula  of  water. 

In  all  reductions  by  hydrogen  peroxide,  each  molecule  of  the  latter  removes  but 
one  atomic  weight  of  oxygen.  Whether  it  behaves  thus  because  its  two  hydro- 
gen units  combine  with  this  oxygen  and  all  its  own  oxygen  escapes,  or  because 
it  furnishes  water  and  one  oxygen  unit  of  the  pair  required  to  form  the  mole- 
cule of  free  oxygen  (the  substance  reduced  furnishing  the  other) ,  has  not  been 
determined. 

The  above  action  is  used  in  quantitative  analysis  for  estimating 
the  quantity  of  hydrogen  peroxide  in  a  given  liquid  after  the  liquid 
has  been  acidified.  The  amount  of  a  standard  (p.  236)  solution  of  the 
permanganate  which  is  required  to  decompose  all  the  peroxide  is 
measured  by  means  of  a  burette  (q.v.").  The  permanganate  is  deep 
reddish-purple  in  color,  while  the  products  are  colorless.  Hence,  after 
the  peroxide  is  exhausted,  the  next  drop  of  the  permanganate  confers 
a  distinct,  permanent,  pink  tinge  upon  the  liquid.  The  addition  of 
the  permanganate  solution  is  stopped  so  soon  as  this  condition  is 
reached  and  the  volume  of  it  that  has  been  used  is  read  off. 

Thermochemistry  of  Hydrogen  Feroxide»  —  The  formation  of 
hydrogen  peroxide  from  the  free  elements  is  accompanied  by  evolution 
of  heat : 

H,  +  02  =  H202  Aq  +  45,300  cal. 

Hence  the  substance  is  formed  by  direct  union  (p.  304).  But  its  de- 
composition into  water  and  oxygen  gives  out  a  further  supply  of 
heat  : 

H2  02  =  H20  +  0  +  23,100  cal. 

The  sum  of  these  two  stages,  of  course,  yields  the  same  result 
(cf.  p.  78)  as  the  direct  formation  of  water  (68,400  cal.). 

When  hydrogen  peroxide  is  used,  instead  of  free  oxygen,  for  oxidiz- 
ing purposes,  each  such  action  liberates  23,100  calories  of  heat  more 
in  the  former  case  than  it  would  in  the  latter.  Hence  the  activity  of 
the  substance  as  an  oxidizer  (cf.  p.  271). 


308  INORGANIC   CHEMISTRY 

Peroxides  :  Chemical  Constitution  and  Molecular  Structure. 

—  To   represent    the  chemical  behavior   of  hydrogen  peroxide,   two 
different  graphic  formulae  (cf.  p.  279)  have  been  proposed  : 


H-0  H 

| 
H-0 


and        X0  =  O. 


In  the  latter  of  the  two  formulae,  one  of  the  oxygen  units  holds  four 
equivalents  of  other  materials  instead  of  two.  Oxygen  being  or- 
dinarily bivalent,  the  two  extra  valences  may  plausibly  be  supposed 
to  involve  a  feebler  state  of  combination,  and  therefore  to  portray  the 
tendency  of  the  compound  easily  to  give  up  one  unit  of  oxygen. 
The  linking  of  the  oxygen  units,  common  to  both  formulae,  expresses 
the  fact  that  we  do  not  obtain  hydrogen  peroxide  from  substances  con- 
taining less  than  two  units  of  this  element  in  each  mole.  Thus,  to  be 
consistent,  we  write  Na^Og,  although  we  have  no  means  of  determin- 
ing the  molecular  weight  of  this  particular  peroxide.  The  former  of 
the  above  formulae  is  more  generally  used. 

All  oxides  containing  two  units  of  oxygen  do  not  yield  hydrogen 
peroxide,  however.  Thus,  lead  dioxide  (Pb0.2)  gives,  by  interaction 
with  dilute  acids,  water  and  oxygen  only.  Carrying  out  our  system, 
therefore,  we  make  barium  and  hydrogen  peroxides  alike,  and  we 
assign  different  constitutions  (p.  224)  to  barium  and  lead  dioxides  and 
different  structures  to  their  molecules  : 


0        HC1  Cl        H-O 

|    +  ->Ba'        +  | 

0         HC1  XC1         H-0 


O  Cl 

Pbf     +2HCl-»Pbx       +  H20+0. 
^O  XC1 

In  confirmation  of  this  we  find  that  lead  can  form  an  unstable  tetra- 
chloride  and  other  compounds  in  which,  as  in  PbIVO2,  it  is  quadriva- 
lent. But  barium  gives  no  other  compounds  in  which  there  is  even 
the  semblance  of  quadrivalence.  So  our  structural  formula  ingeni- 
ously leaves  it  bivalent  even  in  Ba02.  We  assign,  therefore,  to  all 
substances  which  give  hydrogen  peroxide,  and  are  therefore  true 
peroxides,  the  "  peroxide  structure,"  consisting  of  linked  oxygen  units, 
while  for  those  which  give  no  hydrogen  peroxide,  we  write  formulae  in 
which  the  oxygen  units  are  independent  of  one  another. 


OZONE   AND   HYDROGEN   PEROXIDE  309 

Exercises.  —  1.  What  volume  of  ozone  will  be  taken  up  by  100  c.c. 
of  water  at  12°  from  a  stream  of  oxygen  containing  7.5  per  cent  of 
ozone  (p.  155)  ? 

2.  Formulate   the   action   of   carbon  dioxide   on   barium  dioxide 
(p.  303)  after  the  manner  of  that  of  sulphuric  acid  on  the  same  sub- 
stance (p.  304).     The  dissolving  gas  gives  an  additional  equilibrium  : 
C02  (gas)  +  H2O  <=±  H2C08  (diss'd). 

3.  At  what  temperature  will  a  ten  per  cent  solution  of  hydrogen 
peroxide  freeze  (p.  305)  ? 

4.  Write  the  thermochemical  equations   for   oxidation  of  indigo 
by  ozone  (pp.  271,  302)  and  by  hydrogen  peroxide. 


CHAPTER   XIX 
ELECTROLYSIS 

Introductory.  —  Experiment  shows  that  most  solutions  which 
exhibit  chemical  transformations  by  interchange  of  groups  (p.  283),  and 
all  which  show  evidence  of  dissociation  by  measurements  of  osmotic 
pressure,  freezing-point  depression,  and  boiling-point  elevation,  are 
precisely  those  which  are  conductors  of  electricity  and  suffer  decom- 
position by  the  passage  of  the  electric  current.  Solutions  which, 
on  the  contrary,  behave  normally  in  respect  to  the  three  physical  prop- 
erties are  nonconductors.  Solutions  of  the  former  kind,  consisting  of 
ionogens  (p.  296)  dissolved,  usually,  in  water,  are  called  electrolytes, 
and  the  effect  of  an  electric  current  upon  them  is  named  electrolysis. 
Solutions  of  the  same  substances  in  toluene,  chloroform,  etc.,  do  not 
conduct  electricity,  are  not  decomposed  by  the  current,  and,  as  we  have 
seen  (p.  294),  show  no  evidence  of  ionization.  Substances  like  sugar, 
methyl  acetate,  etc.,  which  show  no  evidence  of  dissociation,  are  non- 
conductors, whatever  solvent  we  employ. 

In  endeavoring  to  connect  these  important  facts,  we  must  remem- 
ber that  the  pure  solvent  by  itself,  and  the  pure  substances  which  we 
dissolve  in  the  solvent  by  themselves  are,  at  ordinary  temperatures,  all 
but  complete  nonconductors  of  electricity.  Dry  salts,  except  when  at 
high  temperatures  and  fused,  and  dry  hydrogen  chloride  (p.  182),  on 
the  one  hand,  and  pure  water  (p.  95)  on  the  other,  do  not  permit  the 
passage  of  the  current.  Yet  a  mixture  of  one  of  the  former  with  the 
latter  conducts  extremely  well.  Let  us  consider  first  the  nature  of 
the  decomposition  which  accompanies  the  conduction. 

Chemical  Changes  Connected  with  Electrolysis.  —  When  the 
wires  from  a  battery  are  attached  to  platinum  plates  immersed  in  any 
electrolyte  (e.g.  Fig.  65,  p.  169),  we  observe  that  the  products  appearing 
at  the  two  electrodes  are  always  different.  They  may  be  of  several 
kinds  physically,  and  will  be  secured  for  examination  variously  ac- 
cording to  their  nature.  When  they  are  gases  which  are  not  too 
soluble,  they  may  be  collected  in  inverted  tubes  filled  with  the  solu- 

310 


ELECTROLYSIS  311 

tion.  Solids,  if  insoluble  in  the  liquid,  will  either  remain  attached  to 
the  electrode  or  fall  to  the  bottom  of  the  vessel  as  precipitates.  Sol- 
uble substances  on  the  other  hand  will  usually  not  be  visible.  They 
may  be  handled  by  interposing  a  porous  partition  of  some  description 
which  will  restrain  the  diffusion  of  the  dissolved  body  away  from  the 
neighborhood  of  the  electrode,  while  not  interfering  appreciably  with 
the  passage  of  the  current.  Surrounding  one  electrode  with  a  porous 
battery  jar  is  a  convenient  method  for  effecting  this. 

When  a  current  of  electricity  is  passed  in  this  fashion  through  a 
solution  of  silver  nitrate  AgN08,  we  observe  that  at  the  negative 
electrode  metallic  silver  is  set  free  and  adheres  to  the  plate.  At  the 
positive  electrode  a  gas,  oxygen,  appears.  Since  these  substances  do 
not  account  for  all  the  constituents  of  the  salt,  we  are  impelled  to  ex- 
amine the  solution  around  each  pole,  and  discover  that  nitric  acid 
(HN08)  is  being  formed  along  with  the  oxygen  at  the  positive  end. 
Although  we  have  now  found  something  more  than  the  two  parts  of 
the  original  molecule,  we  have  little  difficulty  in  explaining  the  pres- 
ence of  these  two  products  on  the  assumption  that  the  original 
molecules  were  divided  into  the  parts,  Ag  and  N08.  The  latter,  since 
it  is  not  a  known  compound,  must  have  interacted  with  the  water  to 
produce  nitric  acid  and  oxygen:  2ISrO8 -f  H2O  — » 2HNO8  +  0.  The 
atomic  oxygen  has  subsequently  united  so  as  to  form  the  gas  (O2). 
The  whole  change  may  therefore  be  tabulated  as  follows : 

Neg.  Wire,  Ag.  < Ag.NO8 >  O2  and  HNOg,  Pos.  Wire. 

If  we  substitute  cupric  nitrate  Cu£N"08)2,  we  obtain  a  red  deposit 
of  metallic  copper  on  the  negative  plate,  and  at  the  positive  plate 
oxygen  and  nitric  acid  are  formed.  We  infer  therefore  that  the  parts 
of  the  original  molecule  are  Cu  and  N08 : 

Neg.  Wire,  Cu  < Cu.(N08)2 >  O2  and  HlSTOg,  Pos.  Wire. 

With  a  solution  of  potassium  nitrate  we  find  hydrogen  and  oxygen 
appearing  at  the  negative  and  positive  electrodes  respectively. 
Litmus  paper,  however,  shows  the  presence  in  the  solution  of  a  base 
(potassium  hydroxide,  KOH)  at  the  negative  and  an  acid  (nitric  acid)  at 
the  positive  end.  Secondary  chemical  changes  have  occurred  at  both 
poles.  We  infer  that  the  parts  of  the  parent  molecules  are  K  and  NOr 
The  former,  in  its  customary  manner  (p.  99),  instead  of  being 
liberated,  gave  rise  to  free  hydrogen  and  potassium  hydroxide, 


312  INORGANIC   CHEMISTRY 

2K  +  2H20  ->  2KOH  +  H2 : 

Neg.  Wire,  H2  and  KOH< K.N08 >O2and  HN03,  Pos.  Wire. 

We  are  confirmed  in  these  conclusions  when  we  employ  a  pool  of  mer- 
cury in  place  of  the  negative  wire.  A  portion  of  the  potassium  is 
found  to  have  dissolved  in  the  mercury  and  escaped  interaction  with 
the  water. 

When  dilute  sulphuric  acid  is  electrolyzed  (p.  95),  the  result  is  a 
liberation  of  hydrogen  and  oxygen  and  an  accumulation  of  sulphuric 
acid  round  the  positive  electrode  : 

Neg.  Wire,  H2<-     -H2.S04—      ->  O2  and  H2S04,  Pos.  Wire. 

All  acids  give  hydrogen  alone  at  the  negative  electrode. 

Of  the  various  illustrations  we  have  encountered,  the  electrolysis 
of  hydrochloric  acid  (p.  184)  happens  to  be  the  only  one  which  delivers 
the  two  components  (H  and  Cl)  with  the  minimum  of  modification  by 
secondary  interaction.  The  gases  liberated  are,  of  course,  H.J  and  Clj. 
The  chlorides,  bromides,  and  iodides  of  all  such  metals  as  do  not 
interact  with  water  give  equally  simple  results.  Secondary  actions  are 
equally  worthy  of  notice,  not  only  because  they  are  common,  but  also 
because  they  often  play  a  part  in  industrial  electrolytic  processes. 

On  now  comparing  the  chemical  behavior  of  a  large  number  of 
ionogens  with  the  results  of  electrolysis  of  the  same  substances,  we  are 
led  to  identical  conclusions.  Acids  contain  hydrogen  possessing 
certain  special  properties  (p.  281),  and  by  electrolysis  they  divide  so  as 
to  give  up  this  constituent  alone  at  one  electrode.  Salts  undergo 
double  decomposition  easily  and  exchange  radicals  with  other  ionogens 
(p.  283),  and  the  current  divides  their  molecules  at  the  same  point, 
liberating  the  radicals. 

Quite  a  crop  of  problems  is  raised  by  this  discovery :  Since  a 
solution  may  eventually  be  cleared  of  all  the  hydrochloric  acid,  for 
example,  which  it  contains,  we  should  like  to  know  how  the  con- 
stituents in  the  center  of  the  cell  reach  the  electrodes.  Then  there  is 
the  question  of  the  quantity  of  electricity  required  to  effect  a  given 
amount  of  decomposition  in  a  given  ionogen.  Finally,  since  the 
details  are,  as  usual,  inscrutable,  a  formulative  hypothesis  will  be 
needed  to  explain  the  whole  proceeding. 

Ionic  Migration.  —  The  first  of  these  questions  is  easily  answered 
by  experiment.  We  have  only  to  take  an  ionogen  one  of  whose  radi- 


ELECTROLYSIS 


313 


cals  is  colored,  and  watch  the  movement  of  the  colored  material  as  it 
drifts  towards  the  electrode.  Thus,  in  dilute  cupric  sulphate  solution, 
a  freezing-point  determination  shows  that  the  depression  has  practi- 
cally double  the  normal  value.  In  other  words,  the  dissociation 

CuS04^Cu"-f  SO/' 

is  almost  complete.*  Now,  the  blue  color  of  this  solution  cannot  be 
due  to  the  remaining  molecules  of  CuS04,  for  anhydrous  cupric  sul- 
phate is  colorless.  Nor  is  it  due  to  the  color  of  the  S04"  ion,  for 
dilute  potassium  sulphate  and  dilute  sulphuric  acid  are  both  color- 
less. On  the  other  hand,  all  cupric  salts,  in  dilute  solution,  have  the 
same  tint.  The  color  is  therefore  that  of  the  cupric  ion  (Cu").  Simi- 
larly the  deep-yellow  tint  of  potassium  dichromate  K2Cr2O7  (q.v.)  is 
that  of  the  Cr2G7"  ion.  Hence  either  of  these  substances  will  serve 
the  purpose  of  showing  how  the  ionic  material  moves. 

One  of  the  above  salts  is  dissolved  in  warm  water  containing  about 
5  per  cent  of  agar-agar,  and  the  lower  part  of  the  U-tube  (Fig.  77)  is 
charged  with  the  mixture.  After 
this  fluid  has  set  to  a  jelly,  a  few 
grains  of  powdered  charcoal  are 
added  on  each  side  to  mark  the 
present  limits  of  the  colored  ions, 
or  strips  of  paper  are  pasted  on 
the  outside.  Then  any  colorless 
electrolyte,  such  as  potassium 
nitrate  solution,  is  added  on  each 
side,  and  the  electrodes  are  hung 
one  in  each  limb.  The  lower  half 
of  the  potassium  nitrate  solution 
on  each  side  contains  agar-agar 
also.  The  agar-agar  does  not  offer 
any  appreciable  resistance  to  the 
motion  of  the  ions,  and  is  pre- 
sumed to  form  a  sort  of  open  network  in  the  solution.  It  is  added  to 
prevent  all  motion  of  the  water.  Immersion  of  the  whole  tube  in  ice 
and  water  prevents  the  melting  of  the  jelly  by  the  heat  generated 
by  the  current. 

*  In  all  ions  the  valence  is  indicated,  as  in  this  equation,  by  the  number  of 
superior  marks  •  or '  as  the  case  may  be.  The  use  of  this  custom  will  appear 
presently. 


FIG.  77. 


314  INORGANIC   CHEMISTRY 

After  a  time,  we  observe  that  the  blue  cupric  ions  ascend  above  the 
mark  on  the  negative  and  descend  away  from  it  on  the  positive  side. 
With  potassium  dichromate  the  yellow  ions  move  in  the  opposite 
direction  with  reference  to  the  poles.  In  each  case  there  is  no  shad- 
ing off  in  the  tint.  The  motion  of  the  whole  aggregate  of  colored  ions 
occurs  in  such  a  way  that,  if  the  contents  of  the  tube  were  not  held  in 
place  by  the  jelly,  we  should  believe  that  a  gradual  motion  of  the 
whole  solution  was  being  observed.  With  a  current  of  110  volts,  and 
a  16-candle  power  lamp  in  series  with  the  cell,  the  effect  becomes 
apparent  in  a  few  minutes. 

Although  the  SO/'  ions  are  invisible,  we  may  safely  infer  that  they 
are  drifting  towards  the  positive  electrode.  Indeed,  this  can  be  demon- 
strated by  interposing  a  shallow  layer  of  jelly  containing  some  barium 
salt  a  little  distance  above  the  charcoal  layer  on  tjie  positive  side. 
When  the  SO/'  ions  reach  this,  barium  sulphate  begins  to  be  precipi- 
tated and  the  layer  becomes  cloudy.  In  a  similar  way  the  progress  of 
other  colorless  ions  may  be  rendered  visible. 

It  appears  therefore  that  electrolysis  is  not  a  local  phenomenon, 
going  on  round  the  electrodes  only,  but  that  the  whole  of  the  dis- 
sociated solute  is  set  in  motion.  It  is  on  account  of  this  remarkable 
property  of  traveling  or  migrating  towards  one  or  other  of  the  elec- 
trodes connected  with  a  battery  that  the  ions  receive  their  name  (Gk. 
tW,  going).  The  term  was  first  applied  by  Faraday  to  the  materials 
liberated  round  the  electrodes. 

Relative  Speed  of  Migration  of  Different  Ions. — The  speeds  of 
different  ions  may  readily  be  compared.  The  cupric  ion  moves  at  the 
same  speed  whatever  salt  of  copper  we  employ.  In  fact,  the  speeds 
of  all  ions  are  individual  properties  and  are  independent  of  the  nature 
of  other  ions  that  may  be  present.  The  speeds  of  all  are  increased 
by  using  a  current  of  greater  electromotive  force.  Under  similar  con- 
ditions, the  relative  speeds  of  most  ions  are  in  the  neighborhood  of 
50  or  60,  on  the  scale  commonly  used  in  expressing  ionic  velocities. 
Thus,  we  have,  K*  65.3,  Cl'  65.9,  Cu"  49.  The  speed  of  the  hydrogen 
ion  is  the  greatest  of  all,  318,  while  that  of  hydroxyl  ion  (OH')  comes 
next,  being  174. 

The  actual  speeds  of  these  ions  in  dilute  solutions  at  18°,  when 
driven  by  a  potential  difference  of  1  volt  between  plates  1  cm. 
apart,  expressed  in  cm.  per  hour  is  :  K*  2.05,  Cl'  2.12,  Cu"  1.6,  H*  10.8, 
OH'  5.6. 


ELECTROLYSIS 


315 


By  an  experiment  similar  to  the  last,  and  devised  by  A.  A.  Noyes,  the  relative 
speeds  of  different  ions  may  be  demonstrated.  The  U-tube  (Fig.  78,  showing  the 
same  tube  A  before  the  current  starts,  and  B  after  it  has  been  passing  for  some 
time)  is  partly  filled  with  agar-agar  emulsion  containing  potassium  chloride  and 
phenolphthalein  (see  Indicators).  On  the  right  side,  a  few  drops  of  potassium 
hydroxide  have  been  added  to  render  the  mixture  pink.  On  the  left,  a  few  drops 
of  hydrochloric  acid  are  present,  and  the  mixture  is  colorless.  Above  the  charcoal 
layer,  in  the  right  limb,  a  mixture  of  hydrochloric  acid  and  cupric  chloride  (i.e., 
H*  and  Cu"),  and  in  the  left  limb  potassium  hydroxide  solution  (i.e.  OH'),  are 
placed.  The  positive  electrode  is  introduced  on  the  right  and  the  negative  on  the 
left.  The  H"  and  Cu"  ions  drift  away  from  the  former  down  the  tube  towards  the 


H*  &  Cu' 


OH'  ^= 


Pink  =  OH 


FIG.  78. 

latter,  the  OH'  ions  away  from  the  latter  down  the  tube  towards  the  former. 
The  motion  of  the  H*  is  marked  by  the  disappearance  of  the  pink  color,  that  of  the 
Cu"  by  the  advance  of  a  blue  layer,  that  of  the  OH'  by  the  progress  of  a  pink 
coloration.*  By  the  time  the  H"  ions  have  been  displaced  5£  cm.,  the  Cu"  ions 
have  moved  1  cm.  and  the  OH'  about  2$  cm.  These  distances  indicate  their  rela- 
tive speeds  of  migration. 

Faraday's  Laws.  —  Having  learned  that  electrolytes  furnish  at 
least  two  definite  decomposition  products  by  the  action  of  the  elec- 
tricity, we  naturally  inquire  next  whether  there  is  any  chemical  relation 
between  the  quantities  of  the  products  set  free  by  the  same  current. 
Quantitative  experiments  in  electrolysis  show  the  most  perfect  adjust- 
ment in  this  respect.  Thus,  in  a  single  cell,  the  quantities  of  material 

*  Bases,  on  account  of  the  OH'  they  give,  turn  phenolphthalein  solutions  from 
colorless  to  pink  ;  acids,  on  account  of  the  H"  they  furnish,  turn  it  from  pink  to 
colorless  (see  Indicators). 


316 


INORGANIC   CHEMISTRY 


liberated  at  the  two  poles  are  invariably  chemical  equivalents  of  one 
another.  With  hydrochloric  acid,  while  1.008  g.  of  hydrogen  is  being 
liberated  at  one  pole,  35.45  g.  of  chlorine  are  set  free  in  the  same  time 
at  the  other.  While  63.6  g.  of  copper  are  being  deposited  from  ctipric 
sulphate  at  one  pole,  96  g.  of  S04  are  being  liberated  at  the  other, 
and,  by  interaction  with  the  water,  form  16  g.  of  oxygen  and  98  g.  of 
sulphuric  acid. 

Again,  the  amount  of  any  one  substance  liberated  is  proportional 
to  the  quantity  of  electricity  which  has  traversed  the  cell.  This  is  the 
first  part  of  Faraday's  law. 

Finally,  the  passage  of  equal  quantities  of  electricity  through  sev- 
eral different  acids  liberates  equal  amounts  of  hydrogen  from  each. 
This  is  true,  whether  the  passage  of  the  given  quantity  of  electricity 
is  compressed  into  a  brief  time  in  one  case  and  spread  over  a  longer 
time  in  another,  or  is  uniform  in  all  cases  compared.  It  is  irrespec- 
tive of  the  state  of  dilution  and  of  the  temperature  of  each  acid. 
Thus  two  moles  of  hydrochloric  acid  are  always  decomposed  for 
every  one  of  sulphuric  acid  by  the  same  current.  Similarly,  if  in  dif- 
ferent cells  we  place  solutions  of  substances  like  sodium  chloride  (JSTaCl), 
cupric  chloride  (CuCl2),  antimony  chloride  (SbCl3),  ferrous  chloride 
(FeCl2),  and  ferric  chloride  (FeCls),  equal  amounts  of  chlorine  are 
liberated  by  currents  of  equal  strength  in  the  same  time  in  each. 

If  we  consider  the  relation  of  these  facts  to  the  equivalence  of  the 
materials  liberated  in  any  one  cell,  it  will  be  evident  that  when  one 
gram  of  hydrogen  is  liberated  from  each  of  the  two  acids  mentioned 

above,  one  equivalent 

HCl         H.SO,       KaCl        CuCl,       SbCl,       Fed,       Fed,   4.  of    Chlorine    will    be  Set 

free  in  the  one  cell,  and 
one  equivalent  of  SO4, 
or  half  the  weight  rep- 
resented by  the  for- 
mula, will  be  set  free 
in  the  other.  Similarly, 
with  the  chlorides  of 

the  first  three  metals,  while  35.45  g.  of  chlorine  are  being  liberated  in 
each  cell,  the  quantities  of  the  metal  set  free  will  be,  of  sodium  23  g., 
of  copper  one-half  of  63.6  g.,  and  of  antimony  one-third  of  122  g. 
Finally,  with  the  two  iron  salts,  the  quantities  of  iron  liberated  by  the 
same  current  will  be  one-half  and  one-third  of  56  grams  respectively. 
The  simplest  way  in  which  to  insure  the  passage  of  precise^  equal 


FIG.  79. 


ELECTROLYSIS  317 

amounts  of  electricity  through  all  the  cells  is  to  arrange  them  in 
series.  We  know  that  in  such  circumstances  the  quantity  of  electri- 
city traversing  any  section  of  the  whole  circuit  must  be  the  same  as  that 
traversing  any  other.  In  a  series  of  cells  containing  substances  like 
the  above,  therefore,  during  the  time  that  1.008  g.  of  hydrogen  is  being 
set  free,  we  shall  have  liberation  of  the  equivalent  quantities  of  each 
of  the  other  ions  (Fig.  79).  Thus  the  second  part  of  Faraday's  law 
states  that  chemically  equivalent  quantities  of  ions  are  liberated  by 
the  passage  of  equal  quantities  of  electricity. 

The  Ionic  Hypothesis.  —  The  main  facts  in  regard  to  electrolytic 
conduction  being  now  before  us,  the  problem  is  to  adapt  the  atomic  and 
molecular  hypotheses  to  their  explanation,  that  is,  to  their  detailed  de- 
scription. There  are  two  peculiarities  to  be  accounted  for  : 

How  can  the  production  of  a  conducting  medium  by  mixing  two 
nonconductors  be  imagined  to  take  place  ?  The  solvent  is  a  non- 
conductor, and  the  ions,  even  if  they  are  composed  of  conducting  mate- 
rial, are  distributed  through  the  liquid  as  independent  particles  and 
cannot  furnish  a  continuous  medium  for  the  stream  of  electricity. 
This  will  be  clear  when  we  remember  that  although  liquid  mercury  is 
an  excellent  conductor,  mercury  vapor,  composed  as  it  is  of  conduct- 
ing particles,  is  not  a  conductor. 

Again,  the  conducting  power  of  the  solution  is  indissolubly  con- 
nected with  the  fact  that  the  original  molecules  of  the  solute  have  been 
broken  up  by  the  solvent  into  smaller  molecules  containing  one  or 
more  atoms.  Why  should  this  particular  kind  of  sub-molecules  be  at- 
tracted by  electrically  charged  plates,  which  have  been  lowered  into 
the  solution,  when  molecules  of  dissolved  sugar,  for  example,  are  not 
so  attracted? 

An  answer  to  the  second  question  readily  suggests  itself.  The 
only  bodies  which  we  find  to  be  conspicuously  attracted  by  electrically 
charged  objects  are  bodies  which  are  already  provided  with  electric 
charges  of  their  own.  Thus  we  are  led  to  add  to  the  molecular  hypoth- 
esis the  assumption  that  substances  which  undergo  dissociation  in 
solution  divide  themselves  into  a  special  kind  of  electrically  charged 
molecules.  Since  the  solution,  as  a  whole,  has  itself  no  charge,  equal 
quantities  of  positive  and  negative  electricity  must  be  produced  : 

HC1  <=>  H  +  01     NaCl  <=±  Na  +  Cl    NaOH  <=>  Na  +  OH. 
Wild  as  this  supposition  seems  at  first  sight  to  be,  it  turns  out  that 


318  INORGANIC   CHEMISTRY 

no  valid  objection  to  it  can  be  raised.     That  it  furnishes  an  answer  to 
both  of  our  questions  must  first  be  shown. 

A  battery  is  a  machine  which  maintains  two  points,  its  poles,  or 
two  wires  connected  with  them,  at  a  constant  difference  of  potential. 
One  cell  of  a  storage  battery,  for  example,  maintains  a  potential  dif- 
ference of  two  volts.  When  the  wires  are  joined,  directly  or  indi- 
rectly, the  poles  are  immediately  discharged,  but  the  cell  continuously 
reproduces  the  difference  in  potential  by  generating  fresh  electricity. 
Now  the  effect  of  immersing  two  plates,  one  of  which  is  kept  by  the 
battery  at  a  definite  positive  potential  and  the  other  at  a  definite  nega- 
tive potential,  into  a  liquid  filled  with  floating  multitudes  of  minute 
bodies,  already  highly  charged,  may  easily  be  foreseen. 

The  figure   (Fig.  80)  will 

Cathode  +  Anode 

cation  =  Ag  convey  some  idea  of  the  be- 

anion  =  NO,— >        + J 

havior  of  the  parts  of  a  system 
such  as  we  have  imagined. 
The  electrodes  are  marked  — 
and  -f.  The  negatively 
charged  plate  attracts  all  the 


•  e 

e  "  ©  ® 


... 

9  ©     •   •  e  e 

Q 


e 


*        e  e  e    e         O 

6  e —  ^      positively  charged  particles  in 

I'm  the  vessel,  and,  although  these 


so.  particles  are  in  continuous  and 

irregular  motion,  they  never- 
theless begin,  on  the  whole,  to  drift  toward  the  plate  in  question.  On  the 
other  hand,  the  negatively  charged  particles  are  repelled  by  this  plate 
and  attracted  by  the  positive  plate,  so  that  they  drift  in  the  opposite 
direction.  Those  which  are  nearest  each  plate,  on  coining  in  contact  with 
it,  will  lose  their  charges  of  electricity,  turning  thereby  into  the  ordinary 
free  forms  of  the  matter  of  which  they  are  composed.  The  continuous 
removal  of  the  electrical  charges  of  the  plates  through  contact  with  ions 
of  the  opposite  charge  furnishes  occasion  for  recharging  of  the  plate  from 
the  battery,  and  thus  gives  rise  to  a  continuous  current  in  each  wire. 
Again,  the  continuous  drifting  of  positively  and  negatively  charged 
particles  in  opposite  directions  through  the  liquid,  constitutes  what,  in 
the  view  of  all  external  means  of  observation,  appears  to  be  an  electri- 
cal current  also.  A  magnetized  needle,  for  example,  which  is  deflected 
when  brought  near  one  of  the  wires  of  the  battery,  is  influenced  in  the 
same  way  by  being  brought  over  the  liquid  between  the  electrodes. 
The  illusion,  so  to  speak,  of  an  electric  current  is  complete,  although 
in  reality  it  is  a  convection  of  electricity  that  is  taking  place.  Further- 


ELECTROLYSIS  319 

more,  the  quantity  of  electricity  being  transported  across  any  section 
of  the  whole  system  is  the  same  as  that  across  any  other,  whether 
this  section  be  taken  through  one  of  the  wires,  through  the  electrolyte, 
or  even  through  the  battery  at  any  point.  As  fast  as  the  ions  are  thus 
annihilated  as  such,  the  undissociated  molecules  are  drawn  upon  for 
the  production  of  fresh  ones,  as  in  all  chemical  equilibria.  Eventually, 
by  continuing  the  process  long  enough,  if  the  substances  set  free  are 
actually  deposited  and  do  not  go  into  solution  again  in  any  form,  the 
liquid  can  be  entirely  deprived  of  the  solute  which  it  contains. 

The  analogy  to  the  transportation  of  a  fluid  like  water  is  noticeable, 
although  not  complete.  Water  may  be  transported  in  three  ways.  It- 
may  flow  through  a  pipe,  it  may  pass  by  pouring  freely  from  one  con- 
tainer to  another,  and  it  may  be  carried  in  vessels.  Thus  a  stream  of 
water,  essentially  continuous,  might  be  arranged,  in  which  part  of  the 
passage  took  place  through  the  pipes,  part  by  pouring  from  the  pipes 
into  buckets,  and  part  by  the  carrying  of  those  buckets  between  the  ends 
of  the  pipes.  The  quantity  of  water  passing  a  given  point  per  minute 
in  this  system  would  be  the  same  at  every  part,  although  the  actual 
method  by  which  the  water  was  transported  past  the  various  points 
might  be  different.  In  such  a  disjointed  circuit  we  suppose  the  elec- 
tricity to  move  when  carried  from  a  battery  through  an  electrolytic 
cell.  It  flows  in  the  wire,  passes  by  discharge  between  the  pole  and 
the  ion,  and  is  transported  upon  the  ions  in  the  liquid.  The  parallel 
is  imperfect,  however,  because  we  have  used  the  conception  of  two 
electric  fluids  and  because  the  ions  are  already  charged  in  the  solution, 
and  before  any  connection  with  the  battery  is  made.  They  do  not,  so  to 
speak,  transport  the  electricity  of  the  battery,  but  their  own. 

Difficulties  Presented  by  this  Hypothesis.  —  The  question  was 
raised  (p.  295),  as  to  how  we  can  imagine  separate  atoms  of  sodium  to 
exist  in  water  without  acting  upon  it,  as  the  metal  sodium  usually  does. 
But  the  ions  of  sodium  in  sodium  chloride  solution  are  not  metallic 
sodium.  They  bear  large  charges  of  electricity.  They  possess  an 
entirely  different,  and  in  fact,  by  measurement,  much  smaller  amount 
of  chemical  energy  than  free  sodium.  And,  as  we  have  seen,  the 
properties  of  a  substance  are  determined  as  much  by  the  energy  it  con- 
tains as  by  the  kind  of  matter.  Metallic  sodium  and  ionic  sodium  are, 
simply,  different  substances. 

We  think  of  hydrogen  chloride  and  common  salt  as  exceedingly 
stable  substances,  and  are  averse  to  believing  that  precisely  these  com- 


320  INORGANIC  CHEMISTRY 

pounds  should  be  highly  dissociated  by  mere  solution  in  water.  But  it 
must  be  remembered  that  in  solution  they  undergo  chemical  change  very 
easily,  and  it  is  only  in  the  dry  form  that  they  show  unusual  stability. 

Again,  why  do  not  the  ions  combine  in  response  to  the  attractions 
of  their  charges  ?  The  answer  is  that  they  do  combine,  but  the  rate 
at  which  combination  takes  place  is  no  greater  than  that  at  which  the 
molecules  decompose,  so  that  on  the  whole  the  proportion  of  ions  to 
molecules  remains  unchanged. 

Finally,  it  might  appear  that  the  assumption  that  bodies  could 
retain  high  charges  in  the  midst  of  water  is  contrary  to  all  experience. 
It  must  be  remembered,  however,  that  the  molecular,  pure  water, 
which  separates  the  ions  from  one  another,  is  a  perfect  nonconductor. 
The  moisture  which  covers  electrical  apparatus  and  causes  leakage  of 
static  electricity  is  not  pure  water,  but  a  dilute  solution  containing 
carbonic  acid  (p.  114)  and  materials  from  the  glass  of  which  the  appa- 
ratus is  made.  It  conducts  away  the  charge  electrolytically,  by  means 
of  the  ions  it  contains,  and  not  by  itself  acting  as  a  conductor. 

Amounts  of  Electricity  on  the  Ions.  —  Since  one  atomic  weight 
of  chlorine  (35.45  parts)  and  one  atomic  weight  of  hydrogen  (1.008 
parts)  are  simultaneously  liberated  in  the  electrolysis  of  hydrochloric 
acid  and  the  solution  remains  electrically  neutral,  it  follows  that  the 
ions  of  hydrogen  and  chlorine  bear  equal  charges.  Again,  since  the 
same  thing  is  true  of  one  atomic  weight  of  copper  (63.6  parts)  and  two 
atomic  weights  of  chlorine  (70.9  parts)  in  cupric  chloride,  it  follows 
that  the  cupric  ion  bears  a  double  charge.  Similarly,  the  antimony 
(p.  316)  ions  bear  triple  charges.  In  ferrous  and  ferric  salts  the 
amounts  carried  by  the  iron  ions  differ,  being  in  the  one  case  twice, 
and  in  the  other  three  times  as  great  as  those  carried  by  the  hydrogen 
or  chlorine  ions.  We  may  sum  this  up,  then,  by  saying  that  univalent 
ions  possess  the  same  quantity  of  electricity,  and  other  ions  bear 
quantities  greater  than  this,  in  proportion  to  their  valence.  This  rule 
is  simply  a  restatement  of  Faraday's  law,  which  it  brings  into  direct 
relation  to  the  ionic  hypothesis. 

In  writing  equations  involving  ions,  the   numbers   of    +   and  — 

charges  must  always  be  equal : 

+        —  +++        — 

KC1  ^  K  +  Cl  SbCl,  <±  Sb  +  3C1 

CuCl2  <=>  Cu  +  2C1  FeCL,  <±  Fe  +  2C1 

++      =  +++        — 

t  Cu  +  S04  FeCls  <=>  Fe  +  3C1 


ELECTROLYSIS  321 

In  order  that  this  idea  may  be  carried  out  consistently,  the  libera- 
tion of  any  of  these  ionic  materials  at  one  electrode  in  electrolysis  is 
written  as  follows  : 

K  +  ©_>K  2C1  4-  2  0  -»  CV 

Here  ©  and  0  represent  the  unit  quantities  of  negative  and  posi- 
tive electricity  furnished  by  the  battery  to  the  electrodes  and  destroyed 
by  opposite  charges  upon  the  ions. 

The  harmony  between  the  quantity  of  electricity  and  the  chemi- 
cal valence  of  the  material  which  it  liberates  is  complete.  The 
picture  which  the  process  of  electrolysis  in  a  series  of  cells  (p.  316) 
presents  to  our  minds  is  very  interesting.  The  progress  of  the  elec- 
tricity through  the  series  is  accompanied  by  a  simultaneous  discharge 
in  all  the  cells  of  chemically  corresponding  numbers  of  atoms.  For 
every  atom  of  antimony  that  is  liberated  in  one  cell,  three  atoms  of 
chlorine,  three  atoms  of  hydrogen,  and  one  atom  of  ferric  iron,  are  set 
free  at  the  same  time.  For  two  atoms  of  ferric  iron,  three  atoms  of 
ferrous  iron  and  three  atoms  of  copper  are  deposited.  Even  in  the 
battery  which  generates  the  current,  the  chemical  changes  taking  place 
proceed  atom  for  atom  and  valence  for  valence  in  unison  with  those  in 
the  cells  on  the  circuit.  For  example,  if  the  battery  contains  zinc  plates, 
for  every  atom  of  zinc  that  dissolves,  one  of  copper  and  two  of  chlorine 
will  be  liberated  in  one  of  the  cells.  Our  imaginary  mechanism  thus 
puts  all  the  processes  going  on  in  the  circuit  in  the  light  of  move- 
ments of  the  parts  of  a  perfectly  adjusted  and  interlocked  machine. 

Resume  and  Nomenclature.  —  An  ion  may  be  denned  as,  an 
atom  or  group  of  atoms  bearing  a  positive  or  negative  charge  of 
electricity,  and  formed  through  the  dissociation  of  an  ionogen  by  a 
solvent  like  water. 

Each  molecule  gives  two  kinds  of  ions  with  opposite  charges. 
These  two  are  forthwith  distinct  and  independent  substances,  save  that 
the  attractions  of  the  charges  prevent  separation  by  diffusion.  They 
differ  from  non-ionic  substances  of  the  same  material  composition  when 
such  are  known.  The  electrical  charge  is  one  of  the  essential  con- 
stituents, and  when  it  is  removed  the  properties  alter  entirely.  Thus  we 
have  two  kinds  of  hydrogen,  gaseous  molecular  hydrogen  (H2),  and  ionic 

hydrogen  (H),  with  entirely  different  chemical  properties  (p.  295). 
Since  the   writing    of   the  +  and  —   charges   over   the    symbols 


322 


INORGANIC  CHEMISTRY 


occupies  much,  space,  we  shall  hereafter  employ  a  dot  for  the  former 
and  a  little  dash  for  the  latter:  H',  Cl',  Cu",  SO/',  Fe",  Fe"*,  NH'4. 
In  the  ions  formed  from  one  molecule,  the  number  of  dots  and  dashes 
must  be  equal. 

Since  ionic  hydrogen,  ionic  chlorine,  etc.,  are  entirely  different  in 
physical  and  chemical  properties  from  the  corresponding  free  ele- 
ments, they  should  receive  separate  names.  When  it  is  inconvenient 
to  say  "  ionic  hydrogen/'  "  nitrate  ions "  (NO3'),  etc.,  the  following, 
based  on  Walker's  system,  will  be  used  : 


SYM- 
BOL. 

NAME. 

ANION  OF 

SYM- 
BOL. 

NAME. 

CATION  OF 
SALTS  OF 

SO/' 

Sulphanion 

Sulphates 

NV 

Natrion 

Sodium 

SO" 

Sulphosion 

Sulphites 

Ca" 

Calcion 

Calcium 

CIO/ 

Perchloranion 

Perchlorates 

Cu" 

Dicuprion 

Cupric  copper 

CIO/ 

Chloranion 

Chlorates 

K* 

Kalion 

Potassium 

el- 

Chloridion 

Chlorides 

Fe— 

Triferrion 

Ferric  iron 

s'' 

Sulphidion 

Sulphides 

NH/ 

Amrnonion 

Ammonium 

NO,' 

Nitranion 

Nitrates 

Fe'1 

Diferrion 

Ferrous  iron 

OH' 

Hydroxidion 

Hydroxides  (bases) 

H* 

Hydrion 

Hydrogen(acids) 

Faraday  distinguished  the  two  kinds  of  material  which  proceed 
with  and  against  the  positive  current  by  name.  His  terminology  is 
still  used.  Ions  which  proceed  in  the  same  direction  as  the  positive 
current  are  called  cations  (G-k.  Kara,  down).  Such  are  H*,  Cu",  K*, 
NH4*.  They  are  metallic  elements,  or  groups  which  play  the  part  of 
a  metal.  The  electrode  (Gk.  oSds,  a  path)  upon  which  they  are 
deposited,  the  negative  electrode,  is  spoken  of  as  the  cathode 
(Gk.  17  Ka0o8os,  the  way  down). 

The  particles  which  move  in  the  direction  of  the  negative  current, 
and  against  that  of  the  positive,  are  named  anions  (Gk.  avd,  up).  The 
ions  Or,  N08',  SO/',  MnO/  are  of  this  kind.  They  are  usually  com- 
posed of  non-metals,  although  sometimes,  as  in  MnO/,  the  components 
may  be  partially  metallic.  They  are  set  free  at  the  positive  electrode, 
which  is  therefore  named  the  anode  (Gk.  17  a»/o8os,  the  way  up). 
Chemists  speak  of  metals  and  non-metals  as  positive  and  negative 
elements,  respectively  (cf.  p.  119),  even  when  electrical  relations  are 
not  directly  in  question,  and  ions  are  not  concerned. 


Actual  Quantities  of  Electricity  Concerned  in  Electrolysis.— 

The  coulomb  is  the  unit  quantity  of  electricity  :  the  liberation  of  one 


ELECTROLYSIS  323 

gram  of  hydrogen  corresponds  to  the  passage  of  96,540  coulombs 
round  the  circuit.  In  other  words,  one  gram  of  hydrion  (H*)  carries 
this  quantity  of  electricity.  Equivalent  amounts  of  other  ions,  for 
example,  108  g.  of  argention  (Ag')  and  -9^-  g.  of  sulphanion  (SO/'), 
carry  the  same  quantity.  Thus  the  unit  quantity  of  electricity,  one 
coulomb,  deposits,  or  is  carried  by  ^si?tf  £•  °f  hydrogen  or  ^£§f  ^  g. 
of  silver. 

In  consequence  of  this,  the  deposition  of  silver  or  copper  is  used 
as  a  means  of  measuring  quantities  of  electricity.  The  increase  in 
weight  of  the  negative  electrode  in  a  cell,  called  under  such  circum- 
stances a  voltameter,  is  a  measure  of  the  quantity  of  electricity  which 
passes  around  the  whole  circuit  of  which  it  forms  a  part. 

It  is  more  common  to  define  the  quantity  of  electricity  in  terms  of 
current  strength.  A  current  of  such  nature  that  one  coulomb  flows 
through  the  system  per  second  is  said  to  have  a  strength  of  one 
ampere.  Such  a  current  liberates  ^£¥¥  g.  of  hydrogen  or  s£$f  IF  g.  of 
silver  per  second.  A  current  which  deposits  twice  as  much,  has  two 
amperes  current  strength. 

From  this  it  is  evident  that  a  current  of  one  ampere  would  take 
96,540  seconds,  or  twenty-six  hours  and  49  minutes,  to  deposit  1  g.  of 
hydrogen  (about  11  liters),  or  108  g.  of  silver,  or  48  g.  of  sulphanion. 
A  current  of  five  amperes  would  accomplish  the  same  result  in  a  fifth 
of  the  time. 

The  Electrical  Energy  Required  to  Decompose  Different 
Compounds.  —  Chemical  compounds  are  of  very  different  degrees  of 
stability,  and  hence  the  quantities  of  energy,  electrical  or  otherwise, 
required  to  decompose  them  vary  widely.  Thus,  hydrogen  chloride  is 
very  stable,  while  hydrogen  iodide  is  easily  decomposed  by  heating. 
The  disunion  of  equivalent  quantities  of  these  substances  in  aqueous 
solution  absorbs  39,300  cal.  and  13,100  cal.  of  heat  energy,  respectively. 
Hence,  although  equal  quantities  of  electricity  (96,540  coulombs  in 
each  case)  perform  this  office,  very  unequal  amounts  of  electrical  energy 
are  used  up  in  the  electrolysis. 

The  energy  in  a  stream  of  water  is  represented  by  the  product  of 
the  quantity  passing  a  given  section  and  the  pressure  or  head  of  water 
available  at  that  point.  If  the  pressure  is  low,  the  work  that  can  be 
done  will  be  small,  even  if  the  quantity  flowing  is  great.  So  electrical 
energy  is  expressed  by  the  product  of  the  current  strength,  or  quantity 
passing  per  second  during  a  certain  period  of  time,  and  the  electro- 


824  INORGANIC   CHEMISTRY 

motive  force.  The  latter  corresponds  to  pressure,  and  is  defined  by 
the  difference  in  potential  of  two  points  in  the  circuit  between  which 
the  energy  is  being  used  up. 

Now,  in  the  series  of  cells  which  was  described  (p.  316),  each  cell, 
while  being  traversed  by  the  same  quantity  of  electricity  as  any  of 
the  others,  cuts  down  the  electromotive  force  of  the  current  in  propor- 
tion to  the  amount  of  energy  consumed  by  the  decomposition  going  on 
within  it.  Hence,  while  a  voltmeter  will  show  no  difference  in  poten- 
tial between  two  neighboring  parts  of  the  heavy  wires  used  as  connec- 
tions, for  no  work  is  being  done  in  the  wires,  it  will  show  a  considerable 
difference  in  potential  between  two  points  which  are  separated  by  one 
of  the  cells. 

A  system  of  cars  hauled  by  a  cable  is  analogous  to  our  set  of  cells 
and  more  familiar.  When  clutched  to  the  cable,  all  the  cars  move 
with  equal  speed,  but,  being  loaded  with  different  numbers  of  passen- 
gers, take  very  different  amounts  of  power  from  the  moving  cable. 

We  should  infer  from  this,  that  to  decompose  every  electrolyte,  a 
current  of  a  certain  minimum  electromotive  force,  sufficient  to  furnish 
the  fall  in  potential  necessitated  by  the  chemical  change,  which  would 
be  different  in  different  cases,  would  be  required.  This  is  found  to  be 
the  case.  For  the  easy  decomposition  of  sulphuric  acid  and  liberation 
of  the  products  an  electromotive  force  of  at  least  1.92  volts  is  necessary, 
for  hydrochloric  acid  1.41  volts,  for  hydriodic  acid  0.53  volts,  for  zinc 
sulphate  2.7  volts,  and  for  silver  nitrate  0.70  volts.  When  we  use  a 
current  of  electromotive  force  falling  short  of  that  specified,  we  find 
that  the  flow  of  electricity  is  interrupted.  The  electrolytic  cell  practi- 
cally forms  a  break  in  the  circuit,  (see  Chap,  xxxviii). 

Polarization.  —  It  is  found  that  when  plates  of  platinum,  a  metal 
which  is  not  acted  upon  by  the  liberated  radicals,  are  used,  the  products 
of  electrolysis  accumulate  on  the  electrodes  and  tend  to  produce  a  re- 
verse current  (see  Electromotive  chemistry).  The  cell  is  said  to  be 
polarized.  Thus,  after  hydrochloric  acid  has  been  electrolyzed  for  a 
few  moments,  hydrogen  and  chlorine  adhering  to  the  two  platinum 
plates  set  up  this  current.*  If  the  battery  is  disconnected,  the  elec- 
trolytic cell  becomes  for  a  brief  time  itself  a  battery,  the  re-ionization 
of  the  hydrogen  and  chlorine  (reproducing  hydrochloric  acid)  furnish- 
ing the  energy.  It  is  the  continuous  overcoming  of  this  reverse  current, 

*  If  copper  plates  are  used,  cupric  chloride  is  formed  at  the  positive  plate 
(anode),  and  no  polarization  can  occur  at  that  plate. 


ELECTROLYSIS  325 

and  prevention  of  the  reunion,  that  demands  the  minimum  electro- 
motive force  (here  1.41  volts)  of  which  mention  has  just  been  made. 

It  is  possible  to  arrange  cells  in  which  no  polarization  can  take 
place.  Thus,  when  we  electrolyze  cuprie  sulphate  between  copper 
electrodes,  the  copper  is  deposited  upon  one  plate  and  the  S04  removes 
the  copper  from  the  other  plate,  forming  cuprie  sulphate,  thus  restoring 
the  electrolyte  to  its  original  condition.  The  only  difference  is  that  a 
portion  of  the  copper  has  been  deposited  on  one  pole  and  an  equivalent 
amount  has  been  removed  from  the  other  (see  Copper  refining).  With 
such  cells,  no  minimum  difference  in  potential  is  required  to  effect  elec- 
trolysis, for  there  is  no  polarization  current  to  be  overcome.  The 
feeblest  electric  current  will  produce  continuous,  if  slow,  chemical 
change. 

This  result  is  extremely  interesting,  for  it  shows  that  the  operation 
of  electrolysis  in  itself  does  not  require  the  consumption  of  much 
energy.  If  the  molecules  were  actually  torn  apart  by  the  electricity, 
then  all  electrolytic  operations  would  require  a  minimum  electromotive 
force  for  their  maintenance.  The  fact  just  stated,  therefore,  is  signifi- 
cant, for  it  confirms  the  present  views  in  regard  to  the  theory  of  solu- 
tions. It  is  in  agreement  with  the  belief  that  the  actual  production  of 
the  ions  is  accomplished  by  the  water  in  advance,  and  quite  independ- 
ently of  the  use  of  electricity,  and  that  the  sole  function  of  the  elec- 
tricity in  the  process  of  electrolysis  within  the  solution  consists  in  the 
pilotage  of  the  ions  in  reverse  directions  according  to  their  charges,  an 
operation  which  necessarily  consumes  but  little  energy.  The  friction 
alone  of  the  moving  ions  has  to  be  overcome.  It  makes  clear  the  fact 
that  it  is  only  when  the  chemical  change  in  the  cell  involves  the  actual 
decomposition  of  some  material,  accompanied  (as  in  the  electrolysis  of 
hydrochloric  acid)  by  the  final  delivery  of  the  constituents  in  the  free 
state,  that  considerable  consumption  of  electrical  energy,  proportional 
to  the  extent  of  the  chemical  change,  must  take  place. 

Conductivity  for  Electricity.  —  The  facility  with  which  equi- 
molar  solutions  of  different  substances  conduct  electricity,  when  they 
are  placed  under  like  conditions,  depends  jointly  on  the  degree  of 
ionization,  on  the  speed  with  which  the  ions  move,  and  on  the  valence 
of  the  ions.  When  equivalent  instead  of  equimolar  amounts  are  com- 
pared, the  last  of  these  factors  drops  out  of  consideration.  The  most 
highly  dissociated  acids,  as  we  should  expect,  since  they  give  large 
numbers  of  the  speedy  hydrogen  ions,  are  the  best  conductors.  The 


326  INORGANIC   CHEMISTRY 

highly  ionized  bases,  such  as  potassium  and  sodium  hydroxide,  come 
next.  The  best  conductors  among  salts  fall  considerably  behind  both 
of  these,  because,  although  their  degrees  of  ionization  may  not  be  less 
than  those  of  the  best  conducting  acids  and  bases,  their  ions  all  move 
more  slowly  than  do  hydrion  and  hydroxidion.  On  the  other  hand, 
concentrated  solutions  all  conduct  badly,  relatively  to  the  number  of 
molecules  originally  used  in  making  them,  because  only  that  propor- 
tion of  the  substance  which  is  ionized  contributes  to  the  conduction 
(p.  318).  All  this  is  just  what  we  should  expect,  in  view  of  our 
hypothesis,  for  the  passage  of  the  electricity  must  be  dependent  upon 
the  frequency  with  which  discharges  of  the  ions  upon  the  electrodes 
occur,  and  this,  in  turn,  must  depend  upon  both  the  concentration  and 
the  speed  of  the  ions.  To  return  to  an  analogy  used  once  before,  the 
rate  at  which  a  fluid  can  be  transferred  between  two  reservoirs  must 
depend  upon  the  denseness  of  the  array  of  buckets  available  and  on 
the  speed  with  which  they  are  moved  and  on  their  individual  capacity. 

Ordinarily,  it  is  the  resistance  which  a  substance  presents  to  the 
passage  of  the  electric  current  which  is  measured.  Obviously,  how- 
ever, for  the  present  purpose  it  is  more  convenient  to  give  expression 
to  the  reciprocal  of  this  value,  which  we  term  the  conductivity.  In 
order  that  the  results  may  have  chemical  significance,  we  express  them 
in  terms  of  the  conducting  power  of  one  gram-equivalent  of  the  com- 
pound dissolved  in  water  and  placed  in  a  narrow  cell  whose  opposite 
walls,  of  great  area  and  situated  one  centimeter  apart,  form  the 
electrodes.  Since  the  water  is  a  nonconductor,  the  conducting  power 
of  the  solution  intervening  between  the  plates  is  a  measure  of  the 
capacity  of  the  dissolved  substance  for  facilitating  the  discharge 
between  the  poles.  Inasmuch  as  varying  the  amount  of  the  solvent 
will  not  affect  the  velocity  of  the  ions  of  the  substance,  any  alteration 
in  conducting  power  resulting  from  dilution  must  depend  solely  on 
the  change  in  the  number  of  ions  available  for  carrying  off  the  elec- 
tricity. The  conductivities  of  solutions  of  the  same  substance  in 
different  concentrations  must  therefore  be  proportional  to  the  degrees 
of  its  ionization. 

A  trough  and  amperemeter  *  (Fig.  81)  may  be  used  to  illustrate 
this  principle.  The  electrodes  are  here  long  strips  of  copper  foil, 
which  pass  down  at  the  ends  of  the  trough  and  are  situated,  not  one 
centimeter,  but  ten  or  fifteen  centimeters  apart,  in  order  that  the  con- 

*  For  these  experiments  an  amperemeter  of  low  resistance,  0.5-1  ohm,  must  be 
used,  and  a  battery  of  one  or  two  accumulator  cells  is  sufficient. 


ELECTROLYSIS 


327 


tents  of  the  vessel  may  be  more  easily  seeii.  When  the  two  instru- 
ments are  placed  in  circuit  with  a  battery  and  very  pure  water  is 
poured  into  the  cell,  the  amperemeter  does  not  indicate  the  passage 
of  any  current  of  electricity.  When  a  shallow  layer  of  concentrated 
hydrochloric  acid  is  substituted,  the  situation  is  that  a  definite  amount 
of  hydrogen  chloride  dissolved  in  a  small  amount  of  water  forms  one 
of  the  links  in  the  electric  circuit.  The  deflection  of  the  needle  in  the 
amperemeter  indicates  that  a  certain  current  of  electricity  is  able  to 
pass  through  this  acid.  When  now  distilled  water  is  gradually  mixed 
with  the  acid,  the  amount  of  conducting  material,  in  this  case  hydrogen 


FIG.  81. 

chloride,  is  not  altered.  If,  therefore,  its  capacity  for  conducting 
were  not  affected  by  the  dilution,  the  deflection  of  the  needle  in  the 
amperemeter  would  not  change.  The  wider  dissemination  of  the  hydro- 
chloric acid,  tending  to  diminish  the  specific  conductivity  of  the  solu- 
tion, should  be  exactly  compensated  by  the  greater  area  of  the 
electrodes  coming  into  service.  What  is  actually  observed,  however, 
is  a  very  marked  improvement  in  the  conducting  power  of  the  solu- 
tion.* At  first  the  reading  of  the  amperemeter  increases  very  rapidly. 

*  The  interpretation  of  the  experiment  becomes  easier  if  the  trough  is  first 
nearly  filled  with  distilled  water  and  the  absence  of  deflection  noted.  Then  a  layer 
of  concentrated  hydrochloric  acid  is  introduced  below  the  water,  by  means  of  a 
long-stemmed  dropping  funnel,  and  the  deflection  is  read.  Finally,  the  layers  are 
destroyed  by  stirring,  and  a  great  increase  in  the  reading  of  the  amperemeter  ob- 
served. Here  all  the  hydrogen  chloride  and  water  are  between  the  electrodes  at 
the  time  of  the  second  reading,  and  the  greater  value  of  the  third  reading  cannot 
be  attributed  to  the  use  of  a  larger  area  of  the  poles,  but  solely  to  the  redistribution 
of  the  acid  throughout  a  greater  volume. 


328 


INORGANIC   CHEMISTRY 


Later,  the  effect  of  additional  dilution  becomes  less  marked,  until 
finally  it  becomes  very  slight  indeed. 

When  a  saturated  solution  of  cupric  chloride  is  substituted,  dilu- 
tion is  accompanied  by  a  similar  improvement  in  conductivity.  Here 
we  notice,  besides,  that  the  yellowish-green  liquid  with  which  we  start 
changes  to  a  pale  blue,  as  the  molecules  of  cupric  chloride  are  dis- 
sociated and  the  color  of  the  solution  becomes  more  exclusively  that 
of  the  copper  ions.  When  the  solution  has  become  perfectly  blue, 
further  dilution  is  seen  to  affect  the  conductivity  but  slightly. 

The  approach  to  a  maximum  of  conductivity  reached  in  these  two 
cases  indicates  that  practically  the  whole  of  the  material  has  assumed 
the  ionic  form.  Theoretically  the  absolute  maximum  would  be  reached 
at  infinite  dilution.  The  conductivity  of  the  same  amount  of  sub- 
stance in  more  limited  dilution  is  that  of  the  proportion  of  ions  corre- 
sponding to  this  dilution,  since  the  complete  molecules,  still  present, 
are  without  influence  on  conductivity.  Thus  the  ratio  of  the  conduc- 
tivity at  a'  given  dilution  to  the  maximum  conductivity  is  equal  to  the 
proportion  of  the  whole  material  ionized  at  the  given  dilution.  From 
a  series  of  measurements  for  a  fixed  amount  of  a  substance  at  different 
dilutions,  after  the  results  have  been  plotted,  we  can  usually  (see, 
however,  below)  ascertain  the  limiting,  maximum  conductivity  by 
graphic  extrapolation.  If  A,,  is  the  conductivity  of  an  equivalent  of 
the  substance  dissolved  in  v  liters  of  water,  and  A.^  the  conductivity 
of  the  same  amount  at  infinite  dilution,  then  A^/A.^  is  the  proportion  of 
molecules  completely  ionized  in  the  former  solution.  Xv  is  called 
the  equivalent  conductivity  at  the  dilution  v. 

The  following  numbers  show  the  equivalent  conductivities  at  18° 
of  solutions  of  four  different  substances,  expressed  in  the  units  always 
employed  for  the  purpose  (which  are  reciprocal  ohms).  The  symbols 
AOJ,  meaning  1  equivalent  in  0.1  1. ;  \v  meaning  1  equivalent  in  1  1. ; 
and  so  forth,  denote  the  concentrations. 


\» 

\ 

\o 

**ioo 

(Gale.) 

Hydrochloric  acid      .     . 
Sodium  chloride    .     .     . 
Sodium  acetate      .    .     . 
Acetic  acid  

64.4 
0.05 

301.0 
74.4 
41.2 
1.32 

351.0 
92.5 
61.1 
4.6 

370.0 
103.0 
70.2 
14.3 

384 
110 
78 
(362) 

It  will  be  seen  from  inspection  of  these  figures  that  the  conductiv- 
ity does  not  improve  much  in  the  case  of  the  first  three  substances 


ELECTROLYSIS  329 

when  a  solution  containing  one  equivalent  in  10 1.  is  diluted  ten  times, 
and  that  further  dilutions,  no  matter  how  extensive,  produce  a  still 
smaller  effect.  On  the  other  hand,  acetic  acid  conducts  very  badly  in 
concentrated  solution,  and,  while  the  conductivity  improves  with  dilu- 
tion, it  is  not  possible  experimentally  to  observe  any  approach  to  the 
maximum  conductivity.  The  conductivity,  in  cases  like  this,  is  still 
far  removed  from  the  maximum  at  dilutions  at  which,  with  other  sub- 
stances, the  maximum  is  nearly  attained. 

In  cases  like  that  of  acetic  acid,  the  conductivity  at  infinite  dilu- 
tion cannot  be  estimated  by  extrapolation.  But  fortunately  another 
method  is  available.  The  values  384,  110,  and  78  for  A.^  in  the  cases 
of  the  first  three  substances  can  be  reached  by  extrapolation,  and  rep- 
resent the  conducting  powers  of  equal  numbers  of  ions,  for  there  are 
equal  numbers  of  equivalents  present  and  no  molecules  remain  un-ion- 
ized.  These  values  are  unequal  solely  because  of  the  differing  speeds 
of  the  ions  concerned.  Each  of  them  derives  its  value  from  numbers 
representing  the  relative  speeds  of  the  two  ions  present,  and  must  be 
the  sum  of  these  two  numbers.  If,  therefore,  we  measure  the  relative 
speeds  of  the  two  ions  (p.  315),  we  can  divide  the  value  of  A^  in  this 
proportion  and  learn  the  part  which  each  ion  contributes  to  the  total. 
Dividing  384  in  this  way  we  get  the  speed  of  H*  =  318  and  of  Cl'  = 
65.9  already  given  (p.  314).  Dividing  A.^  for  sodium  acetate  (78),  simi- 
larly, we  get  the  speeds  Na*=  44.4  and  C2H302'  =  33.7.  The  speeds  of 
Cl'  and  Na*  together  (110.3)  must  then  equal  A.^  for  NaCl,  and,  as  we 
see,  they  do.  Similarly,  the  speeds  of  H"  and  C2H802'  together  (351.7) 
must  equal  A.^  for  HC2H302',  although  we  cannot  observe  the  latter 
directly.  This  method  can  be  applied  to  all  of  the  less  highly  ionized 
acids  and  bases,  for  their  sodium  and  potassium  salts  belong  invariably 
to  the  class  of  substances  which  are  most  ionized,  and  for  which,  there- 
fore, A^  can  be  determined  accurately  by  extrapolation. 

Degrees  of  lonization  of  Common  Substances.  —  The  rule, 
degree  of  ionization  =  A^/A.^  (p.  328),  enables  us  to  calculate  the 
value  for  any  dilution,  when  the  necessary  data  are  given.  We  need 
only  the  values  of  the  conductivity  (A^, )  for  different  dilutions  (p.  328) 
and  those  of  the  relative  speeds  of  each  kind  of  ions  expressed  in  the 
same  units.  The  latter,  when  added,  give  Aw. 

Thus,  hydrogen  chloride  in  a  solution  containing  1  equivalent  in 
0.1 1.  (365  g.  per  liter),  which  would  be  a  rather  concentrated  hydro- 
chloric acid,  shows  the  degree  of  ionization  ~^  >  or  0.168  (=  16.8  per 


330 


INORGANIC   CHEMISTRY 


cent).     Normal  hydrochloric  acid  is  ionized  to  the  extent  of  |^>  or 


0.784  ;    normal  sodium  chloride,  7~  >  or  0.676  ;    normal   acetic   acid, 

1  32 

3^2  >  0.004  (=  0.4  per  cent). 

Misapprehension  easily  arises  in  regard  to  the  inferences  that  may  be  drawn 
from  a  conductivity  value.  A  single  such  value,  say  that  for  salt  at  10  1.  dilution 
(92.5),  gives  no  information  about  the  extent  of  ionization.  We  must  have  the 
value  at  infinite  dilution  as  well,  that  is,  we  must  have  the  other  term  of  the  ratio 
corresponding  to  complete  ionization,  before  the  proportion  of  the  molecules 
ionized  at  the  10  1.  dilution  can  be  known.  Further,  we  must  have  the  values  of 
both  for  the  same  salt,  at  the  same  temperature  and  in  the  same  solvent,  for  the 
values  at  all  dilutions  change  markedly  when  any  one  of  these  conditions  is  altered. 
Thus  the  conductivity  of  normal  sodium  chloride  solution  at  60°  is  120,  and  is 
therefore  actually  greater  than  that  at  18°  when  the  dilution  is  infinite.  But 
at  50°  the  conductivity  at  infinite  dilution  is  185,  so  that  at  this  temperature  the 
degree  of  ionization  is  3~f$  or  0.65,  about  the  same  as  at  18°.  On  the  other  hand, 
when  a  little  alcohol  is  added  to  the  aqueous  solution,  the  conductivities  all  dimin- 
ish. But  that  at  infinite  dilution  diminishes  also,  so  that  the  proportion  of  the 
material  ionized  does  not  seem  to  be  greatly  affected.  The  chief  effect  of  raising 
the  temperature  is  simply  to  diminish  the  friction  opposing  the  motion  of  the  ions 
and,  therefore,  to  increase  the  conductivity.  The  change  is  about  2  per  cent  for 
each  degree.  Addition  of  alcohol,  on  the  other  hand,  increases  the  friction  and 
diminishes  the  conductivity.  There  is,  however,  a  real,  though  usually  smaller, 
change  in  the  degree  of  ionization  with  change  in  temperature.  When  the  tem- 
perature is  raised,  the  fraction  ionized  increases  or  diminishes  according  as  the 
heat  of  ionization  is  negative  or  positive  (cf.  p.  260),  and  conversely  when  the 
temperature  is  lowered. 

The  following  tables  include  the  common  reagents  and  give  the 
proportions  of  ionized  molecules  (total  molecules  =  1).  Except  where 
otherwise  specified,  the  data  are  for  normal  solutions  at  18°.  In  the 
case  of  acids  containing  more  than  one  displaceable  hydrogen  unit, 
the  kind  of  ionization  on  which  the  figure  is  based  is  indicated  by  a 
period.  Thus  H.HCOS  means  that  the  whole  of  the  ionization  is 
assumed  to  be  into  H*  and  HCO/. 


FRACTION   IONIZED. 
ACIDS. 


Nitric  acid 0.820 

Nitric  acid  (cone.,  62%)  .  .0.096 
Hydrochloric  acid  .  .  .  .0.784 
Hydrochloric  acid  (cone.,  35%)  0.136 
Sulphuric  acid,  H.H.S04  .  .0.510 
Sulphuric  acid  (cone.,  95%)  .0.007 
Hydrofluoric  acid  .  .  .  .  0 . 070 
Oxalic  acid,  H.HC2O4  (N/10, 
25°) 


Tartaric    acid,    H.HT    (N/10, 

25°) 

Acetic  acid 

Acetic  acid  (N/10)  .... 


0.500 

0.082 
0.004 
0.013 


Carbonic  acid,  H.HCO3  (N/10)  0.0017 
Carbonic  acid,  H.HCO3  (N/25)  0.0021 
Hydrogen  sulphide,H.HS (N/10)  0 . 0007 


Boric  acid,  H.H2B03  (N/10) 
Hydrocyanic  acid  (N/10)  . 
Permanganic  acid  (N/2,  25°)  . 
Hydriodic  acid  (N/2,  25°)  .  . 
Hydrobromic  acid  (N/2,  25°)  . 
Perchloric  acid  (N/2,  25°)  .  . 
Chloric  acid  (N/2,  25°)  .  .  . 
Hydrochloric  acid  (N/2,  26°)  . 
Phosphoric  acid,  H.H2PO4  (N/2, 
25°)  


0.0001 

0.0001 

0.933 

0.901 

0.899 

0.880 

0.878 

0.876 

0.170 


ELECTROLYSIS 


331 


BASES. 


Potassium  hydroxide 
Sodium  hydroxide    . 
Barium  hydroxide  . 
Lithium  hydroxide  . 
Ammonium  hydroxide 
Tetramethylammonium 
droxide  (N/16,  25°)  . 


hy- 


0.77 
0.73 
0.69 
0.63 
0.004 

0.96 


Strontium    hydroxide    (N/64, 

26°)         0.93 

Barium  hydroxide  (N/64,  25°)  .0.92 

Calcium      hydroxide      (N  /  64, 
25°) 0.90 

Silver      hydroxide      (N/1783, 
25°) 0.39 


SALTS. 


Sodium  phosphate,  Na2.HPO4 


(N/32) 

0  83 

Cupric  nitrate  (N/16)    . 
Potassium  chlorate  (N/2) 
Sodium  tartrate  (N/32,  25C 
Potassium  chloride 

.  0.80 
.0.79 
)    -(0.78) 
.0.75 

Ammonium  chloride     . 

.  0.74 

Sodium  chloride 

.  0.676 

Sodium  chloride  (N/2)  . 
Sodium  chloride  (N/10) 
Potassium  nitrate    . 
Potassium  acetate    . 

.  0.734 
.  0.839 
.  0.64 
.  0.64 

Calcium  sulphate  (N/100)    .      .0.63 

Silver  nitrate 0 . 58 

Potassium  sulphate   .      .      .      .  0 . 53 
Sodium  acetate     .      .      .      .      .0.53 
Sodium  bicarbonate,    Na.HC03  (0.52) 
Potassium  carbonate       .      .      .  (0.49) 

Sodium  sulphate 0.445 

Zinc  sulphate 0 . 24 

Zinc  chloride 0.48 

Cupric  sulphate 0 . 22 

Mercuric  chloride       .      .      .   «0.01) 
Mercuric  cyanide       ....  Minute 

Degree  of  Ionization  of  Water.  —  If  we  consider  a  liter  of  water 
as  a  normal  solution  in  which  18  g.  (one  mole)  represents  the  solute 
and  the  rest  stands  for  the  solvent,  the  conductivity  for  complete 
ionization  into  H*  and  OH'  would  be  318  +  174  =  492.  The  actual 
ionization  is  one  ten-millionth  part  of  this.  In  other  words,  there  is 
only  one  ten-millionth  of  1  g.  of  hydrion  and  the  same  fraction  of  17  g. 
of  hydroxidion  in  a  liter  of  water.  A  column  of  water  1  cm.  long  con- 
ducts less  well  than  a  column  of  mercury  of  equal  cross-section  and 
over  660,000  miles  in  length. 

General  Remarks  on  these  Values.  —  It  will  be  seen  from  in- 
spection of  the  above  numbers  for  acids  that  the  proportion  of  the 
molecules  ionized  in  solutions  of  equivalent  concentration  varies 
enormously.  Koughly,  the  acids  might  be  divided  into  four  groups : 
those  in  which  the  ionization  exceeds  70  per  cent  in  normal  solutions ; 
those  in  which  it  lies  between  70  and  10  per  cent ;  those  in  which 
it  lies  between  10  and  1  per  cent ;  and  those  in  which  it  is  smaller 
than  1  per  cent.  To  the  first  class  belong  the  acids  which  we 
generally  recognize  as  the  most  active  in  all  their  chemical  relations, 
namely,  nitric  acid,  the  halogen  hydrides,  and  one  or  two  others.  To 
the  second  class  belong  sulphuric  acid  and  phosphoric  acid,  and  they 
are  less  active.  Amongst  the  acids  whose  ionization  lies  between  one 
and  ten  per  cent  are  such  as  hydrofluoric  acid  and  acetic  acid,  and 


332  INOKGANIC   CHEMISTRY 

chemically  they  are  weak.  Carbonic  acid  and  boric  acid  are  of  the 
fourth  class,  and  are  feeble  acids. 

The  bases,  although  less  numerous,  show  that  a  similar  division 
might  be  made,  although  in  the  above  list  only  two  classes  are 
represented,  —  the  strong  bases,  beginning  with  potassium  hydroxide, 
and  the  feeble  bases,  represented  by  ammonium  hydroxide. 

The  salts  show  a  much  greater  uniformity ;  and,  if  the  list  had  been 
extended  so  as  to  include  the  hundreds  of  common  salts  in  constant 
use,  the  vast  majority  would  have  been  found  to  show  degrees  of 
ionization  lying  between  50  and  80  per  cent.  Only  a  few  fall  below 
these  limits.  The  salts  of  mercury  are  almost  the  only  ones  which 
would  belong  to  the  class  of  least  ionized  substances.  Salts,  like  zinc 
sulphate  and  cupric  sulphate,  in  which  both  ions  are  multivalent,  are 
always  much  less  highly  ionized  than  are  salts  (e.g.  ZnCl2)  made  up  of 
either  of  the  same  ions  along  with  a  univalent  ion.  Salts,  however, 
are  ahnost  never  restricted  in  their  degree  of  ionization,  to  an  extent 
sufficient  to  produce  any  noticeable  effect  on  their  chemical  properties 
(see,  however,  Cadmium  iodide  and  Mercuric  cyanide). 

The  relation  between  degree  of  ionization  and  prominence  of  acid 
or  basic  chemical  properties  will  be  developed  in  the  next  chapter. 

Comparison  with  the  Results  Obtained  by  Other  Methods.  — 

The  value  for  the  degree  of  ionization  as  measured  by  the  conductivity 
method  is  coincident  with  that  found  for  the  same  solution  by  a  study 
of  the  abnormalities  in  freezing-  and  boiling-points  and  in  osmotic 
pressure  (Chap.  xvii).  The  electrical  method  gives  accurate  results 
more  easily  than  do  the  others,  however,  and  is  therefore  the  one  most 
frequently  used.  It  was  Svante  Arrhenius,  a  Swedish  chemist,  who, 
in  1887,  first  noted  the  coincidence  in  the  values  and  devised  the  ionic 
hypothesis  to  account  for  it.  From  the  appearance  of  his  remarkable 
memoir  we  date  the  great  development  which  the  study  of  solutions  * 
has  undergone  in  recent  years. 

Exercises.  —  1.  Name  (p.  322)  the  ionic  materials  furnished  by  the 
dissociation  of  potassium  bromate,  silver  bromide,  sodium  periodate 
(NaI04),  permanganic  acid. 

2.  Give  lists  of  other  anions  and  cations  which  have  been  en- 
countered. 

*  The  Scientific  Memoirs,  No.  IV.  (American  Book  Company),  is  a  reprint 
of  the  fundamental  papers  by  Raoult,  van  't  Hoff,  and  Arrhenius. 


ELECTROLYSIS  333 

3.  How  many  coulombs  are  carried  by  and  will  deposit :   20  g.  of 
silver,    15   g.   of  antimony,  30   g.   of    chlorine,   60   g.    phosphanion 

(P04)  ? 

4.  What   current   strength  (in  amperes)   is  required   to   deposit: 
20  g.  of  silver  in  an  hour,  100  g.  of  iodine  in  5  minutes,  60  g.  of  anti- 
mony in  3  hours  ? 

5.  What  is  the  percentage  of  molecules  ionized  in:    deci-normal 
(N/10)  sodium  chloride,  centi-normal  (N/100)  acetic  acid,  centi-normal 
hydrochloric  acid  (p.  328)  ? 

6.  Give  an  experimental  definition  of  the  term  ion.     That  in  the 
text  (p.  321)  is  in  terms  of  the  hypothesis. 


CHAPTER   XX 
THE   CHEMICAL  BEHAVIOR   OF  IONIC  SUBSTANCES 

BEFORE  considering  the  typical  interactions  of  ionogens  in  solution, 
we  must  have  a  clear  conception  of  the  peculiarities  of  these  bodies 
which  are  likely  to  affect  their  behavior.  The  facts  on  which  such 
a  conception  must  be  based  have  been  given  in  preceding  chapters,  and 
all  that  is  now  necessary  is  to  collect  and  apply  these  facts.  On 
account  of  the  coherence  which  they  give  to  the  subject,  the  figures 
of  speech  of  the  ionic  hypothesis  will  be  largely  employed.  The 
reader  will,  therefore,  do  well  to  exercise  especial  care  to  distinguish 
fact  from  fiction. 

In  this  discussion  it  must  be  made  clear  that  aqueous  solutions 
of  ionogens  are  mixtures  containing  several  solutes.  It  must  also 
be  shown  that  each  kind  of  ions  is  a  distinct  substance  with  indi- 
vidual physical  and  chemical  properties.  Next,  salts  being  used  for 
illustration,  the  commonest  kind  of  interaction,  double  decomposition 
between  ionogens,  will  be  discussed.  In  this  connection  precipitation 
brings  up  the  peculiar  state  of  equilibrium  between  the  undissolved 
solute  and  the  complex  of  molecules  and  ions  in  solution.  Applica- 
tion of  the  same  principle  to  special  cases,  such  as  those  of  acids  and 
bases,  then  follows. 

The  discussion  of  systems  in  equilibrium  in  the  present  chapter 
will  be  purely  qualitative.  The  quantitative  consideration  of  ionic 
equilibria  (cf.  p.  297)  is  postponed  until  the  study  of  the  metals  and 
their  compounds  is  taken  up  (see  Chap,  xxxiv). 

Solutions  of  Ionogens  are  Mixtures.  —  We  are  accustomed  to 
regard  a  bottle  of  sodium  chloride  solution  as  containing  but  one  thing, 
aside  from  the  water.  We  must  now  think  of  it  as  containing 
at  least  three  dissolved  substances,  any  one  of  which  might  be  alone 
responsible  for  some  property  of  the  solution.  The  same  idea  must 
accompany  our  use  of  every  solution  of  an  ionogen.  Thus,  in  ordi- 
nary experiments,  in  which  solutions  of  concentration  not  far  from 
normal  are  commonly  used,  we  have  something  like  the  following 
proportions  (p.  330)  of  the  three  main  components  in  six  typical  cases : 

334 


THE   CHEMICAL   BEHAVIOR   OF  IONIC   SUBSTANCES         335 


(68%),  (78%)CuSO4<=±Cu"  +  SO/'(22%), 
(22%)HC1  ^H'+Cl'   (78%),  (99.6  %)HC2H802^±H' 
(23%)KOH<=iK*-hOH'(77%), 


In  solutions  made  from  salts,  the  greater  part,  and  by  far  the  most 
active  part,  of  the  contents  is  almost  always  ionic.  Cupric  sulphate, 
being  a  salt  both  of  whose  ions  are  bivalent  (p.  332),  practically 
illustrates  the  lower  limit  as  regards  quantity  of  ions,  for  only  a  few 
salts  of  mercury  and  cadmium  fall  far  below  it.  The  acids  and  bases 
have  a  wider  range,  and  a  larger  proportion  of  them  are  like  acetic 
acid  and  ammonium  hydroxide  respectively.  Still,  even  when  small 
in  amount,  the  ions  of  acids  and  bases  are  almost  always  much  more 
active  than  the  molecules. 

The  presence  of  still  other  components  in  solutions  of  salts,  arising 
from  interaction  with  the  water,  will  be  noted  later. 

While  most  of  these  facts  are  ascertained  by  more  or  less  remote 
inference  from  physical  and  chemical  properties  of  the  solutions, 
some  of  them  are  evident  to  the  eye  in  certain  cases.  Thus  the  prog- 
ress of  the  ionization  of  a  salt  may  be  seen  if  one  of  the  ions  is 
different  in  color  from  the  molecules.  Cupric  bromide  in  the  solid 
form  is  a  jet  black,  shining,  crystalline  substance.  When  treated 
with  a  small  amount  of  water  it  forms  a  solution  which  is  of  a  deep 
reddish-brown  tint,  giving  no  hint  of  resemblance  to  a  solution  of  any 
cupric  salt.  This  doubtless  represents  the  color  of  the  molecules. 
When  more  water  is  added,  the  deep  brown  gives  place  gradually  to 
green,  and  finally  to  blue.  The  latter  is  the  color  of  the  cuprion 
(Cu**),  and  is  familiar  in  all  solutions  of  cupric  salts.  The  colorless 
nature  of  solutions  of  potassium  and  sodium  bromides  shows  that 
bromidion  (Br')  is  without  color.  Hence,  in  the  present  instance  it  is 
invisible.  We  are  thus  watching  the  progress  of  the  action  : 

CuBr2  <=>  Cu-  +  2Br'. 

If  1  g.  of  the  solid  is  taken,  it  dissolves  in  about  its  own  weight  of 
water,  and  independent  measurement  shows  that  there  is  relatively 
little  ionization.  Hence  the  solution  is  deep  brown.  When  10  c.c.  of 
water  has  been  added,  70  per  cent  of  the  salt  is  ionized,  and  the  solu- 
tion is  green.  With  40  c.c.  of  water,  only  19  per  cent  remains  in 
molecular  form,  and  the  blue  color  of  the  cuprion  entirely  overbears 
the  tint  of  the  molecules.  If,  at  the  green  stage,  we  dissolve  solid 
potassium  bromide  in  the  liquid,  the  high  concentration  of  bromidion 


336  INORGANIC   CHEMISTRY 

which  results  causes  an  extensive  reversal  of  the  dissociation  (cf.  p. 
250),  and  the  molecules,  with  their  brown  color,  become  prominent  again. 
Sufficient  final  dilution  with  water,  however,  reduces  the  concentrations 
of  all  the  ions  once  more,  the  molecules  dissociate,  and  the  brown 
color  is  displaced  by  the  blue  for  the  second  time. 

Each  Kind   of  Ion   in  a  Mixture   Acts   Independently.  — 

Numberless  facts  show  that  each  kind  of  ion,  for  example  cuprion,  has 
an  individual  set  of  physical  and  chemical  properties  and  behaves  in 
many  ways  as  if  alone  present  in  the  solution.  We  shall  meet  with 
much  evidence  of  this  in  the  sequel.  Some  facts  tending  to  prove  it, 
that  have  already  been  given,  may  be  recalled  (cf.  p.  295). 

If,  in  comparing  the  migration  speeds  of  any  element,  say  copper, 
in  different  salts  (p.  314),  they  were  the  motions  of  substances  like 
Cu(N03)2,  CuBr2,  CuS04,  that  we  were  comparing,  all  analogy  teaches 
us  that  the  speeds  with  which  they  would  move  should  vary  widely. 
That  the  blue  color  drifts  always  at  the  same  pace  shows  that  it  is  the 
same  substance,  namely,  cuprion  (Cu"),  that  we  are  observing. 

If,  in  solutions  of  the  different  permanganates,  KMnO4,  NaMn04, 
Ba(Mn04)2,  and  so  forth,  the  dissolved  bodies  were  different  in  each 
case,  we  should  confidently  expect  the  purple  colors  of  the  solutions  to 
differ  markedly  in  shade.  But,  for  dilute  solutions  of  equivalent  con- 
centrations, when  strict  examination  is  made,  the  tints  are  found  to  be 
absolutely  identical.  We  are  therefore  simply  comparing  different 
mixtures  all  containing  the  same  proportion  of  the  same  free,  colored 
body,  MnO/. 

In  phosphorus  pentachloride  vapor  (p.  255),  the  fully  liberated 
trichloride  and  chlorine  are  prominent  components.  Diminishing  the 
volume  of  a  fixed  amount  of  this  mixture,  by  compression,  throws 
more  chlorine  into  combination  and  the  total  absorption  (from  which 
the  greenish-yellow  color  is  derived)  becomes  less,  the  compounds  of 
phosphorus  being  both  colorless.  Increasing  the  volume,  on  the  other 
hand,  promotes  the  dissociation  and  increases  the  total  absorption.  The 
system  of  ions  and  molecules  in  equilibrium  in  a  solution  of  cupric 
bromide,  or  any  other  ionogen,  behaves  in  exactly  the  same  way.  The 
components  possess  and  exhibit  individual  properties,  much  like  the 
components  of  a  gaseous  mixture  (p.  155),  both  in  this  and  in  other 
respects. 

All  solutions  of  acids  are  sour  in  taste,  irrespective  of  the  nature 
of  the  negative  ion,  while  salts  containing  the  same  negative  radical 


THE  CHEMICAL  BEHAVIOR  OF  IONIC   SUBSTANCES         337 

are  not  sour  at  all.  Hence  in  solutions  of  acids  we  are  tasting  the 
same  free  substance,  hydrion  (H*).  Similarly,  in  solutions  of  all 
alkalies,  we  note  the  soapy  taste  of  hydroxidion  (OH'). 

These  illustrations  concern  physical  properties.  In  the  next  sec- 
tion we  shall  learn  that  an  ionic  material,  such  as  bromidion  or  cu- 
prion,  has  specific  chemical  properties  irrespective  of  the  nature  of  its 
concomitants. 

SALTS,  IONIC  DOUBLE   DECOMPOSITION,  PRECIPITATION. 

Salts.  —  We  have  already  seen  that  salts  differ  much  in  solubility 
in  water,  that  some  combine  with  water  to  form  solid  hydrates  (p.  120), 
that  they  interact  with  acids,  bases,  and  other  salts  by  exchange  of 
radicals,  reversibly  (p.  281),  and  that  they  are  all  ionogens. 

Both  the  positive  and  negative  ions  of  salts  may  be  simple  or  com- 
posite, Na.Cl,  Na.N08,  NH4.C1,  NH4.N08.  The  elements  which  can 
form  a  simple  positive  ion  are  known  in  chemistry  as  metals  (p.  119, 
and  see  Chaps,  xxiii  and  xxxii).  Non-metals,  like  nitrogen,  may  be 
present  in  a  positive  ion,  as  in  NH4',  but  never  exclusively.  In  other 
words,  we  know  no  such  substances  as  nitrogen  sulphate,  or  carbon 
nitrate.  Metals,  on  the  other  hand,  are  frequently  found  in  the  nega- 
tive ion,  but  never  constitute  it  exclusively.  They  are  then  usually 
associated  with  oxygen,  as  in  MnO/,  and  Cr207".  Some  ionic  mate- 
rials are  colored,  Cu"  blue,  Or*"  reddish  violet,  Co"*  pink,  MnO/ 
purple,  Cr207"  orange,  but  most  of  them  are  colorless,  K',  Na*,  Zn", 
Cl',  I',  NO/.  The  ions  of  salts  do  not  affect  litmus.  They  vary  in 
taste,  some  being  salt,  some  astringent,  some  bitter.  In  dilute  solu- 
tions they  are  almost  always  numerous  in  comparison  with  the  surviv- 
ing molecules.  They  carry  electricity,  but  relatively  less  well  than 
do  hydrion  and  hydroxidion,  on  account  of  their  slower  migration.  All 
the  known  ionic  materials  are  found  in  solutions  of  salts.  The  only 
ions  which  are  not  characteristic  of  salts,  although  sometimes  occur- 
ring in  their  solutions  (  see  Mixed  ionogens ),  are  hydrion  H*,  and 
hydroxidion,  OH'. 

Double  Decomposition  of  Salts  in  Solution.  —  When  we  mix 
sodium  chloride  (NaCl)  and  silver  nitrate  (AgN03),  both  in  solution  in 
water  (H20),  there  would  seem  to  be  many  different  possibilities  of 
union  amongst  the  six  elements  represented  in  the  mixture.  But  only 
one  sort  of  change  occurs,  and  it  takes  place  almost  completely: 

AgN08  +  NaCl  <=>  AgCl{+  NaNO8.  (1) 


338  INORGANIC   CHEMISTRY 

So  in  general,  we  find  that  a  single  crosswise  union  of  ions  with  oppo- 
site electrical  charges  is  by  far  the  commonest  kind  of  interaction 
between  ionogens.  There  are  indeed  four  other  kinds  of  ionic  chemical 
change,  as  will  be  seen  in  the  sequel,  but  for  the  present  we  shall  dis- 
cuss only  the  cases  of  double  decomposition. 

When  solutions  of  two  ionized  substances  are  mixed,  the  first  reflec- 
tion which  occurs  to  us  is  that  each  of  these  has  been  diluted  by  the 
water  in  which  the  other  was  dissolved,  so  that  the  first  effect  will  be 
to  increase  the  degree  of  ionization  of  both  to  a  certain  extent.  The 
next  consideration  is,  however,  that  we  have  produced  a  mixture  of 
four  ions,  which  must  have  at  least  some  tendency  to  unite  crosswise. 
Thus  potassium  chloride  and  sodium  nitrate  in  dilute  solution  are  very 
greatly  ionized  before  mixing.  The  reversible  actions,  represented  by 
the  horizontal  pair  of  the  following  equations,  have  taken  place  exten- 
sively. But,  by  mixing  the  liquids,  we  have  brought  into  presence  of 
one  another  two  new  pairs  of  positive  and  negative  ions.  Hence,  two 
other  reversible  actions,  the  vertical  ones, 

KCI^K-   +  cr 

NaN08«=;N(V  +  Na* 

IT        IT 

KN03       NaCl 

will  be  set  up  and  will  proceed  until  a  fresh  equilibrium  of  all  the  ions 
with  all  four  kinds  of  molecules  has  been  reached.  Examination  of 
the  solutions  of  these  four  salts,  separately,  shows  that  they  are  in  an 
equal  degree  extensively  ionized  in  dilute  solutions,  so  that  in  this  par- 
ticular case  the  whole  quantity  of  molecules  of  all  kinds  will  not  be 
very  great.  That  this  inference  is  correct  is  shown  by  much  independ- 
ent evidence,  of  which  two  samples  may  be  given. 

The  change  of  ionic  materials  into  molecular  and  vice  versa  is  always 
accompanied  by  absorption  or  liberation  of  heat  (p.  330).  Now  it 
was  for  long  a  matter  of  surprise  that  when  dilute  solutions  of  salts, 
as  distinct  from  pairs  which  included  acids  or  bases  (q.v.),  were  mixed, 
no  heat-change  was  observable.  This  fact  was  called  the  thermoneu- 
trality  of  salts,  but  the  reason  for  it  was  unknown.  Since  salts  are  all 
highly  ionized,  the  reason  is  now  apparent.  Similarly,  no  changes  in 
color  or  volume  accompany  the  mixing  of  dilute  solutions  of  salts. 

Again,  any  of  the  means  which  may  be  used  for  measuring  the 
number  of  molecules,  including  ions,  in  a  solution,  may  be  applied  to 
learning  whether  any  appreciable  proportion  of  the  latter  has  disap- 


THE    CHEMICAL  BEHAVIOR   OF   IONIC   SUBSTANCES         339 

peared.  Osmotic  pressure,  freezing-point,  boiling-point,  and  conduc- 
tivity are  all  applicable.  The  last  lends  itself  best  to  the  purpose  of 
demonstration.  The  cell  and  amperemeter  described  in  the  last  chap- 
ter  (Fig.  81)  may  be  employed.  We  place  in  the  cell  a  one-fourth 
normal  (N/4)  solution  of  potassium  chloride  and  introduce  an  equal 
volume  of  N/4  sodium  nitrate,  in  such  a  way  that  it  forms  a  separate 
layer  beneath  the  other  solution.  We  read  the  amperemeter,  mix  the 
liquids  by  stirring,  so  as  to  permit  chemical  interaction  to  take  place, 
and  then  note  the  conductivity  once  more.  There  is  no  observable 
change  in  the  conducting  power,  and,  therefore,  no  appreciable  change 
in  the  condition  of  the  ionic  substances  has  taken  place. 

We  shall  have  occasion  in  a  later  paragraph  to  show  that,  where 
the  ions  unite  crosswise  to  an  appreciable  extent,  an  experiment  of 
this  kind  shows  a  marked  diminution  in  the  conductivity,  correspond- 
ing to  the  amount  of  combination  that  has  occurred.  In  such  cases, 
also,  heat  is  either  liberated  or  absorbed. 

It  thus  appears  that  when  dilute  solutions  of  salts  are  mixed,  and 
there  is  no  visible  evidence  of  chemical  change,  the  little  that  has  taken 
place  may  be  neglected.  Practically,  mixing  of  this  kind  is  a  physical 
operation.  Of  course,  in  the  event  of  one  product  being  precipitated, 
a  diminution  in  the  conductivity,  corresponding  to  the  amount  of  ionic 
material  removed,  will  be  observed. 

In  view  of  the  above  explanation,  the  old  question  of  whether  such  a  solution 
contains  the  first  pair  of  salts,  or  the  second  pair,  represented  in  the  double  decom- 
position, KC1  +  NaN03  +±  KNO3  +  NaCl,  loses  its  whole  point.  The  solution  con- 
tains neither  the  initial  molecular  substances  nor  the  molecular  products,  in 
appreciable  amount. 

Solution  and  Precipitation  of  Salts.  —  In  a  body  which  is  dis- 
solving or  being  precipitated,  we  are  observing  the  progress  of  a  rever- 
sible physical  operation  (p.  153) : 

Molecules  (undiss'd)  <=±  molecules  (diss'd). 

If  the  molecules  are  sparingly  soluble,  the  forward  action  is  feeble, 
while  the  backward  one,  when  we  start  with  the  same  material  in  solu- 
tion, makes  great  progress.  Conversely,  when  the  substance  is  a  sol- 
uble one,  the  forward  action  comes  to  a  standstill  (or  the  reverse  action 
occurs)  only  when  the  concentration  of  dissolved  molecules  has  be- 
come very  large. 

Now  it  will  be  noted  that  this  mechanical  adjustment  concerns 
only  the  molecules,  and  that  the  ions,  if  there  are  any,  are  involved 


340  INORGANIC   CHEMISTRY 

only  indirectly.  The  ions  are  in  equilibrium  with  the  dissolved  mole- 
cules : 

AB  (undiss'd)  +±  AB  (diss'd)  <=>  A*  +  B'. 

Hence,  when  a  substance  dissolves,  it  does  so  in  molecular  form,  and 
ions  are  subsequently  generated  from  some  of  these  molecules  until 
equilibrium  is  reached.  Conversely,  when  molecules  come  out  of  so- 
lution, as  the  result  of  cooling,  for  example,  the  diminished  concentra- 
tion of  the  central  term  of  the  chain  enables  more  ions  progressively 
to  unite  until  the  whole  system  has  adjusted  itself  to  the  new  condi- 
tions. 

Now  this  has  an  important  bearing  on  the  result  of  mixing  dilute 
solutions  of  two  soluble  salts.  We  have  seen  that  the  concentration  of 
the  molecules  of  the  new  pair  of  salts  is  never  large.  Yet  it  may 
easily  be  in  excess  of  the  amount  which  the  water  can  hold  in  solution, 
if  one  of  the  salts  is  of  the  relatively  insoluble  class.  This  occurs, 
for  example,  when  a  chloride  is  mixed  with  a  salt  of  silver.  The  sys- 
tem of  equilibria,  leaving  out  that  of  Na*  and  N08',  and  rearranging 
(cf.  p.  338),  so  as  to  save  space,  appears  as  follows  : 

NaCl     <=>  Na'   +01'  ) 
- 


The  concentration  of  dissolved  silver  chloride  which  the  solid  can 
maintain  in  solution  being  very  minute,  most  of  this  salt  is  at  once  pre- 
cipitated. The  ions  continue  to  unite  because  the  requisite  concen- 
tration of  molecules,  whose  dissociation  should  bring  their  union  to  a 
standstill,  has  not  been  kept  up.  The  new  molecules  are  in  turn  pre- 
cipitated. The  system  reaches  a  stable  condition  only  when  the  con- 
centration of  chloridion  (01')  and  argention  (Ag*)  has  fallen  to  that 
which  can  be  maintained  by  the  mere  trace  of  molecules  which  the  in- 
solubility of  the  substance  permits  to  remain  in  solution.  When  the 
system  of  equilibria  is  examined,  we  see  at  once  what  the  result  must 
be.  The  removal  of  chloridion  and  argention  enables  the  remaining 
molecules  of  sodium  chloride  and  silver  nitrate  to  become  completely 
ionized.  Thus  the  concentration  of  NaCl  and  AgN08,  of  Ag'  and  01' 
and  of  the  dissolved  AgCl,  all  become  practically  zero.  The  system 
finally  contains  only  molecular,  solid  silver  chloride  and  the  three  sub- 
stances, Na*  +  NO/  1=;  NaN08,  in  equilibrium,  of  which  by  far  the 
greater  part  is  the  ionic.  There  is  therefore  in  the  action  much  detail 
which  the  first  equation  (p.  337)  did  not  show. 


THE   CHEMICAL  BEHAVIOR  OF  IONIC   SUBSTANCES          341 

The  above  explanation  will  apply  to  any  case  of  precipitation 
resulting  from  the  interaction  of  ionogens.  If  the  least  soluble  of  the 
four  salts  is  more  soluble  than  silver  chloride,  more  concentrated  solu- 
tions are  required  to  secure  precipitation.  The  interaction  of  hydro- 
gen chloride  and  sodium  hydrogen  sulphate  (p.  179)  is  of  this  nature  : 

TTp]    ,  _   TT»  i       pi/        \ 

NaCl  (diss'd)  <±  NaCl  (solid). 


It  should  be  noted  that,  strictly  speaking,  the  only  interaction  tak- 
ing place  when  the  solutions  are  mixed  is  the  production  of  the  insol- 
uble body.  In  the  case  of  silver  chloride,  (1)  and  (2),  the  largest  part 
of  the  chemical  action  may  be  formulated  thus  : 

Ag'  +  Cl'fcsAgCl.  (3) 

The  chief  change  that  has  as  yet  befallen  the  ions  of  sodium  nitrate  is 
that  they  have  been  transferred  from  two  separate  vessels  into  one. 
Potentially  the  salt  has  been  formed.  But  the  actual  union  of  its  ions 
to  give  the  second  product  in  the  molecular  condition  : 

Na'  +  N08'-»NaN08,  (4) 

comes  about  only  when,  at  some  subsequent  time,  if  at  all,  the  water  is 
evaporated  away. 

Individual,  Specific  Chemical  Properties  of  Each  Ionic 
Material.  —  We  wrote  the  equation  for  the  formation  of  silver  chlo- 
ride (Ag*  -f  Cl'  —  >  AgCl)  as  if  argention  and  chloridion  were  the  only 
substances  concerned  in  the  action.  Further  study  shows  this  to  be 
justifiable.  Thus,  hydrochloric  acid,  cupric  chloride,  and  dozens  of 
other  chlorides  may  be  used  instead  of  sodium  chloride  and  give  silver 
chloride  just  as  readily.  The  natrion  had  nothing  to  do  with  the 
result.  Of  course  we  cannot  get  a  solution  containing  chloridion 
alone.  Like  a  vessel  in  which  to  make  the  experiment,  some  positive 
ion  is  required.  But,  like  the  rest  of  the  apparatus,  this  ion  may  be 
varied  indefinitely,  is  not  altered  in  the  course  of  the  change,  and  may 
therefore  be  dispensed  with  in  the  equation.  The  nitranion  (NO/) 
which  accompanied  the  argention  is  similarly  a  part  of  the  apparatus, 
for  silver  sulphate  solution  works  just  as  well  as  silver  nitrate. 

That  chloridion  is  a  substance  with  specific  chemical  properties,  is 
easily  demonstrated.  It  forms  silver  chloride  whenever  it  encounters 
argention.  Other  substances,  even  when  they  contain  chlorine,  lack 


342  INOKGANIC  CHEMISTRY 


this  property.  Chloroform  (CHCy  and  chlorobenzene  (C6H6C1),  in  a 
solvent  in  which  ionogens  are  dissociated,  do  not  interact  when  silver 
nitrate  is  added.  They  give  no  chloridion,  and,  in  fact,  remain  un-ion- 
ized.  Potassium  chlorate  (KC103)  and  perchlorate  (KC104)  and  chlor- 
acetic  acid  (HC2H2C102),  with  argention,  fail  likewise  to  give  silver 
chloride.  They  are  ionized,  but  chloridion  is  not  one  of  the  ions 
of  any  of  them.  The  ions  CIO/,  CIO/,  and  C2H2C10/,  have  properties 
of  their  own,  and  their  compounds  with  argention  are  soluble. 

Other  chemical  properties  of  chloridion  are  :  That  it  unites  also 
with  plumbion  (Pb")  and  monomercurion  (Hg*),  forming  insoluble 
chlorides  (p.  185).  It  is  discharged  and  liberated  as  free  chlorine  by 
fluorine  (p.  241)  : 

2H*  +  2C1'  +  F2  ->  2H'  +  C12  +  2F'. 

Since  the  hydrion  is  not  affected  and  many  chlorides  behave  in  a 
similar  manner,  the  positive  ion  may  be  omitted  : 

2C1'+F2-*C12  +  2F'. 

Finally,  chloridion  has  relatively  little  tendency  to  unite  with  other 
ions,  or,  in  other  words,  the  compounds  of  chloridion  with  most  other 
ions  are  highly  ionized.  Thus  it  combines  with  hydrion  to  the  extent 
of  only  22  per  cent  (p.  330)  in  normal  solution.  In  this  respect  it 
differs  markedly  from  free  chloriue,  just  as  hydrion  differs  from 
hydrogen.  The  free  elements  unite  with  vigor  and  completely.  Hy- 
drogen chloride  is  easy  to  dissociate  into  ions,  but  difficult  to  dissociate 
into  its  constituent  elements.  Nothing  could  show  more  strikingly 
than  this  that  the  ionic  materials  have  chemical  properties  of  their 
own. 

Similarly,  barium  salts  and  ordinary  sulphates  give,  when  mixed,  a 
precipitate  of  barium  sulphate.  Here  we  encounter  a  property  of  ba- 
rion  (Ba")  and  sulphanion  (SO/').  But  potassium  ethyl  sulphate 
(KC2H5S04),  in  spite  of  its  name,  will  not  give  this  reaction  with 
a  barium  salt.  Here  electrolysis  shows  that  sulphanion  is  absent  and 
that  the  negative  ion  is  C2H5SO/. 

In  the  same  way  every  other  ionic  material  may  be  shown  to  be  a 
substance  with  an  individual  set  of  physical  (p.  336)  and  chemical 
properties.  Each  salt,  when  dissolved,  gives  two  kinds  (see,  however, 
below)  of  ionic  materials.  The  solution  is  simply  a  mixture,  and  each 
physical  component  forthwith  behaves  towards  ions  capable  of  uniting 


THE   CHEMICAL   BEHAVIOR   OF   IONIC   SUBSTANCES         343 

with  it,  as  if  it  were  alone.  The  other  materials,  ionic  and  molecular, 
which  are  present,  may  remain  essentially  unaffected  throughout  the 
change. 

Application  in  Chemical  Analysis.  —  Since  the  larger  number 
of  ordinary  chemical  substances  are  ionogens,  and  the  most  rapid  and 
simplest  chemical  changes  take  place  when  they  are  in  solution, 
the  various  reactions  of  their  solutions  are  employed  as  tests  for  the 
substances  in  question.  An  advantage  of  the  use  of  the  solutions  is 
that  they  contain  a  mixture  of  two  independent  materials,  the  anion 
and  the  cation,  and  when  these  have  been  identified  successfully  the 
salt  from  which  they  were  formed  is  known.  The  simplicity  to  which 
chemical  analysis  is  thus  reduced  may  be  seen  when  we  consider  that 
twenty-five  common  metals  with  twenty-five  negative  radicals  might  give 
a  total  of  over  six  hundred  different  salts.  If  the  distinct  properties  of 
each  of  these  had  to  be  considered,  the  identification  of  an  unknown 
substance  would  be  very  difficult.  In  solution,  however,  the  problem 
becomes  much  easier.  Every  solution  made  from  a  single  salt  will  con- 
tain but  two  substances  (in  the  main ;  see,  however,  below),  and  the 
problem  reduces  itself  to  ascertaining  which  two,  out  of  a  total  of 
fifty,  are  present  in  any  particular  case. 

As  an  example  of  the  method,  let  us  suppose  that  we  look  first  for  the  positive 
ion.  Most  systems  of  analysis  begin  by  the  addition  of  a  solution  containing 
chloridion,  generally  dilute  hydrochloric  acid,  to  the  liquid.  If  an  ion  is  present 
which  in  combination  with  chloridion  gives  an  insoluble  compound,  a  precipitate 
will  appear.  Amongst  the  common  positive  ions  but  three  are  of  this  kind,  namely, 
argention,  monornercurion,  and  diplumbion.  So  that  the  precipitate,  if  it  appears, 
is  a  chloride  of  one  of  these  three  metals,  and  the  matter  of  distinguishing  between 
the  three  is  quickly  disposed  of  by  further  examination  of  its  properties.  If  no 
precipitate  comes  out,  then  these  three  metals  are  probably  absent,  and  some  fresh 
ion  capable  of  precipitating  another  set  of  positive  ions  is  introduced  (see  Chap, 
xxxvii).  Thus  by  a  process  of  elimination  we  quickly  find  out  whether  any  metal 
ion  is  present,  and,  if  so,  precisely  which  one  it  is. 

The  language  of  analysis  is  frequently  somewhat  loose.  Thus  we  speak  of  the 
addition  of  a  silver  salt  to  a  solution  as  being  a  "  test  for  chlorine."  As  a  matter  of 
fact,  it  is  not  a  test  for  chlorine.  It  is  not  intended  as  a  test  for  free  chlorine, 
nor  will  it  show  the  presence  of  chlorine  in  many  states  of  combination.  It  is  simply 
a  test  for  ionic  chlorine  (Cl'),  and  cannot  give  us  information  in  regard  to  the  presence 
or  absence  of  any  other  form  of  the  element.  So  the  wet-way  tests  for  **  copper," 
"silver,"  etc.,  so  called,  are  tests  for  the  ionic  forms  of  these  elements,  and  not 
for  the  presence  of  the  element  in  every  form.  Even  the  two  kinds  of  copper  and 
mercury  ions,  Cu",  Cu*,  Hg",  Hg",  must  be  classed  as  distinct  substances.  Thus, 
the  last  is  precipitated  by  chloridion  while  the  second  last  is  not,  mercuric  chloride 
(HgCl,)  being  soluble. 


344  INORGANIC   CHEMISTRY 

Hydrolysis  of  Salts.  —  The  natural  ionization  of  water  is  very 
slight,  but  there  are  cases  in  which  its  effects  become  noticeable,  and 
the  interaction  of  its  ions  with  those  of  dissolved  salts  cannot  be  neg- 
lected. For  example,  an  aqueous  solution  of  pure  cupric  sulphate  is 
always  acid  and  therefore  contains  hydrion  : 


Cupric  hydroxide,  being  a  very  feeble  base,  and  comparable  with  water 
itself  in  the  small  extent  to  which  the  solvent  is  able  to  hold  its  ions 
apart,  is  formed  to  a  small  extent.  The  removal  of  some  hydroxidion 
by  this  means  enables  more  of  the  water  to  dissociate.  This,  in  turn, 
furnishes  the  material  for  the  production  of  more  cupric  hydroxide. 
The  action  does  not  proceed  very  far,  but  it  makes  sufficient  progress 
to  leave  a  perceptible  excess  of  hydrion  in  the  liquid  and  to  give  it, 
therefore,  an  acid  reaction.  The  hydrion  combines  slightly,  but  only 
slightly,  with  the  sulphanion,  for  sulphuric  acid  is  a  highly  ionized 
acid.  This  part  of  the  action  has,  therefore,  been  left  out  of  the  dia- 
gram. The  ordinary  equation  for  this  change  would  be  : 

CuS04  +  H20  fc?  Cu(OH)2  +  H2S04. 

The  hydrolysis  is  much  greater  with  sodium  sulphide  (q.v.)  and  anti- 
mony trichloride  (q.v.). 

Again,  soap  solution  is  always  faintly  alkaline  : 

HC16H81CO,(, 

The  sodium  palmitate  is  highly  ionized,  but  palmitic  acid  (HC16H81C02) 
is  hardly  ionized  at  all.  The  final  result  is  the  production  of  a  recog- 
nizable amount  of  hydroxidion  in  the  solution.  Thus,  a  salt  derived 
from  an  acid  and  a  base  of  very  different  degrees  of  activity,  whether 
it  is  the  base  (as  Cu(OH)2)  or  the  acid  (as  palmitic  acid  or  hydrogen 
sulphide,  q.v.)  which  is  the  weaker  member,  is  likely  to  be  more  or 
less  hydrolyzed  by  water.  In  the  former  case  the  solution  is  acid,  in 
the  latter  basic  in  reaction.  Other  things  being  equal,  salts  containing 
bivalent  or  trivalent  radicals  are  more  noticeably  hydrolyzed  than  are 
those  composed  only  of  univalent  radicals. 

Cases  of  this  kind  being  common,  we  are  thus  compelled  to  enlarge 
our  list  of  possible  components  in  the  solutions  of  any  salt.     In  addi- 


THE   CHEMICAL   BEHAVIOR   OF   IONIC   SUBSTANCES         345 

tion  to  the  molecules  and  ions  of  the  salt,  there  are  present,  water  and 
its  ions,  and  the  molecules  of  the  base  and  acid  formed  by  the  union 
of  the  latter  ions  with  the  former.  There  are  thus  no  less  than  eight 
different  components  in  the  mixture. 

ACIDS  AND  BASES  AND  THEIR  DOUBLE  DECOMPOSITION  WITH  SALTS. 

Hydrogen  Salts.  —  The  substances  of  the  composition  HC1, 
H2S04,  and  so  forth,  are  commonly  called  acids,  and  when  more  con- 
venient we  shall  conform  to  this  usage.  But  it  is  only  when  they 
have  been  dissolved  in  water  or  some  other  ionizing  solvent  that  they 
show  the  properties  characteristic  of  acids.  In  fact,  in  terms  of  the 
ionic  hypothesis,  there  is  only  one  acid,  hydrion  (H"),  although  the 
substances  which  give  it  by  dissociation  are  many.  The  parent  sub- 
stances are  salts  of  hydrogen,'  in  which  the  element  hydrogen  plays 
the  part  of  a  metal. 

The  properties  of  the  hydrogen  salts,  that  is,  of  the  original  iono- 
gens,  are  the  same  as  those  of  any  other  ionogens.  They  are  distin- 
guished from  other  ionogens  by  the  fact  that  their  positive  radical 
is  always  hydrogen  and  that  in  solution  they  yield  hydrion.  Their 
negative  radicals  are  all  different,  Cl,  Br,  I,  C108,  C104,  Br03,  N08,  and  so 
forth.  These  radicals  form  the  negative  ions  in  solutions  of  hydrogen 
salts.  In  such  solutions  all  the  properties  of  these  ions  are  the  same 
as  when  they  are  furnished  by  dissolving  other  salts  containing  the 
same  radicals.  Some  hydrogen  salts  have  oxidizing  powers  like  hypo- 
chlorous  acid  (p.  269).  Usually,  they  exchange  radicals  with  the 
other  ionogens.  They  often  do  this  even  when  dissolved  in  non-dis- 
sociating solvents.  They  frequently  do  it  also  in  the  absence  of  a 
solvent,  especially  when  heated.  They  differ  from  one  another  in  the 
matter  of  solubility  in  water,  some  being  almost  insoluble.  By  solu- 
tion in  water  they  give  acids  of  very  different  degrees  of  activity  (see 
Activity  of  acids,  below). 

Hydrion.  —  Hydrion  is  a  colorless  substance  which  exists  in  water 
and  certain  other  solvents  only.  It  is  always  associated  with  an  equiv- 
alent amount  of  some  negative  ion.  It  is  sour  in  taste,  and  its  pres- 
ence is  recognized  by  the  fact  that  it  turns  blue  litmus  red  and  decolor- 
izes pink  phenolphthalein  (see  Indicators,  below)  solutions.  It  confers 
a  high  conducting  power  upon  solutions  in  which  it  is  contained,  on 
account  of  its  great  speed  of  migration.  It  is  univalent  and  combines 


346  INORGANIC  CHEMISTRY 

with  negative  ions,  such  as  hydroxyl  and  the  negative  radicals  of  salts. 
It  is  displaced  by  inetals  like  magnesium.  In  all  these  respects  it  differs 
markedly  from  free  hydrogen  gas. 

The  displacement  of  hydrogen  from  dilute  acids  (p.  95)  now  ap- 
pears in  a  new  light.  The  action  will  be  formulated  thus  : 

Zn  +  2H'  +  SO/'  -»  Zn"  +  SO/'  +  H2. 

The  sulphanion  (SO/') ,  although  zinc  sulphate  is  somewhat  less  ion- 
ized than  sulphuric  acid,  is  not  much  affected  by  the  change  and  may 
be  omitted: 

Zn  +  2H'-»Zn"+H2. 

Thus,  this  action,  which  takes  place  in  the  same  fashion  with  most 
acids,  is  seen  to  be  independent  of  the  .nature  of  the  negative  ion.  It 
consists  simply  in  the  transference  of  the  electric  charges  from  the 
hydrogen  to  the  zinc,  whereby  the  latter  becomes  ionic.  The  dis- 
charged hydrogen  is  liberated  as  gas.  When  the  solution  is  evapo- 
rated, the  ionogen,  in  the  above  case  zinc  sulphate,  is  formed : 

Zn"+  SO/'->ZnS04. 

Modes  of  lonization  of  Acids.  —  An  acid  containing  but  one 
unit  of  hydrogen  in  its  molecule  can  give  but  two  kinds  of  ions. 
Thus,  chloric  acid  gives  only  H*  and  C108'.  When  more  than  one 
hydrogen  unit  is  present,  however,  more  than  two  kinds  of  ions  are 
formed.  Thus,  sulphuric  acid,  H2S04,  produces,  in  the  first  place,  hydro- 
sulphanion : 

H2S04  <±  H-  +  HSO/. 

The  latter  is  also  an  acid,  but  is  considerably  less  active  than  sulphuric 
acid.  Hence,  the  further  dissociation  of  this  ion  (HSO/<=±  H*-f-  SO/') 
lags  considerably  behind  the  primary  dissociation.  In  concentrated 
solutions  of  the  acid  there  is,  therefore,  much  HSO/  present.  In  very 
dilute  solutions,  however,  SO/'  predominates.  We  know  that  HSO/ 
is  a  weaker  acid,  and  is  dissociated  with  greater  difficulty  by  water, 
because  acid  salts  (see  below),  like  KHS04,  which  give  this  ion,  are 
much  weaker  acids  than  are  acids  like  HC1  and  HC108,  with  which 
the  substance  HSO/  might  fairly  be  compared.  This  behavior  is  not 
peculiar  to  sulphuric  acid,  but  is  shown  by  all  acids  containing  more 
than  one  hydrogen  unit  in  the  molecule  (cf.  Hydrogen  sulphide). 


THE   CHEMICAL   BEHAVIOR   OF   IONIC    SUBSTANCES          347 

Activity  of  Acids.  —  In  solutions  containing  equivalent  quantities 
of  hydrogen  salts,  and  therefore  equal  amounts  of  combined  hydrogen, 
in  equal  volumes,  the  concentration  of  hydrion  present  at  any  moment 
in  each  will  be  different.  This  concentration  will  be  high  or  low 
according  to  the  extent  to  which  water  is  able  to  dissociate  the  mole- 
cules. Now  the  activity  of  the  hydrion,  that  is,  the  speed  with  which 
it  will  interact,  like  that  of  any  other  substance,  depends  on  its  con- 
centration (p.  250).  Hence  the  hydrogen  salts  furnish,  on  being  dis- 
solved, acids  of  all  degrees  of  activity.  Thus  in  normal  hydrochloric 
acid,  the  fraction  dissociated  is  0.78,  and  the  hydrion  is  0.78-normal, 
whereas  in  normal  acetic  acid  the  hydrion  is  only  0.004-normal 
(p.  330).  Yet  the  amounts  of  hydrogen  chloride  and  hydrogen  acetate 
per  liter  contain  equal  quantities  of  combined  hydrogen,  namely,  1  g. 
each.  Both  the  solutions  in  fact  are  normal  in  respect  to  combined 
hydrogen.  But  the  normal  acetic  acid  has  only  about  one  two-hun- 
dredth of  the  activity  of  normal  hydrochloric  acid. 

That  a  difference  in  the  activity  of  different  acids  does  exist  may 
be  shown,  roughly,  by  placing  similar  pieces  of  the  same  metal,  say  zinc, 
in  equal  volumes  of  various  normal  solutions  of  acids,  such  as 
hydrochloric,  sulphuric,  and  acetic.  The  hydrogen  is  evolved  more 
rapidly  by  the  first  than  by  the  second,  and  very  much  faster  by 
either  than  by  the  last.  Naturally,  the  first  is  sooner  exhausted,  while 
the  third  acts  in  its  slow  way  for  a  very  long  time  before  being  all 
used  up.  In  the  third  case  few  ions  of  hydrogen  are  at  hand  at  any 
one  moment,  but  more  are  formed  continuously  from  the  molecules,  to 
take  the  place  of  those  displaced.  Thus  the  total  amount  of  hydrogen 
obtained  from  each  acid  is  finally  the  same.  It  is  the  speed  of  evolu- 
tion alone  which  is  different  and  shows  the  differing  concentrations  of 
the  hydrion. 

In  cases  of  extremely  small  ionization,  the  presence  or  absence  of 
visible  action  on  litmus  may  form  another  means  of  estimating 
activity.  Thus,  litmus  is  easily  turned  red  by  a  deci-normal  solution 
of  acetic  acid  or  of  any  more  active  acid  (p.  330),  but  hydrogen  sul- 
phide, in  a  solution  of  the  same  molecular  concentration,  contains  only 
one-twentieth  as  many  hydrogen  ions  (p.  330),  and  affects  litmus  paper 
but  slightly.  Paper  dipped  in  Congo  red  exhibits  differences  in 
the  activity  of  acids  by  the  different  depths  of  the  tints  it  assumes. 
For  example,  it  is  much  less  markedly  affected  by  acetic  than  by  sul- 
phuric acid  of  the  same  concentration  (see  Indicators,  below). 

Many  hydrogen  salts  are  but  slightly  soluble.     Thus,  with  silicic 


348  INORGANIC   CHEMISTRY 

acid  (q.v.),  the  solid  can  keep  only  a  small  concentration  of  molecules 
in  solution:  H2Si03  (solid)  <=?  H2Si08  (diss'd).  So  that,  although  some 
ions  are  doubtless  present,  H2Si03  (diss'd)  ^=>  2H*  -f  Si03",  their  con- 
centration, being  dependent  on  that  of  the  molecules,  is  very  minute 
indeed.  Still,  even  in  the  absence  of  an  effect  upon  litmus,  the  sub- 
stance can  be  recognized  to  be  an  acid.  Thus,  by  the  action  of  sodium 
hydroxide,  silicic  acid  can  be  made  into  sodium  silicate,  NagSiOg,  which 
is  highly  soluble  and  highly  ionized.  Hence,  since  Si03"  is  a  negative 
ion,  we  reach  the  conclusion  indirectly  that  H2Si03  is  an  acid. 

Substances  like  ammonia  NH3,  sugar,  and  alcohol,  although  they 
contain  hydrogen,  are  not  hydrogen  salts.  They  are  not  ionogens 
(cf.  p.  281),  and  give  no  hydrion.  lonizable  and  non-ionizable  hydrogen 
may  even  be  contained  in  the  same  compound.  Thus,  each  molecule  of 
acetic  acid  (HC2H302)  contains  four  hydrogen  units,  but  gives  only  one 
hydrogen  ion.  The  other  three  are  part  of  the  acetanion  (C2H302'). 
We  infer  this  because  metals  can  be  substituted  for  one  hydrogen  unit 
(NaC2H302),  but  not  more. 

Salts  of  Hydroocyl.  —  Substances  like  potassium  hydroxide,  am- 
monium hydroxide  (NH4OH),  and  zinc  hydroxide  (Zn(OH)2),  are 
commonly  called  bases.  But  it  is  only  in  their  aqueous  solutions  that 
the  basic  properties  appear.  There  is  only  one  base,  namely,  hydroxid- 
ion  (OH'),  and  these  substances  are  simply  the  source  of  it.  The  parent 
substances  are  salts  of  some  metal,  or  group  playing  the  part  of  a 
metal  (e.g.  NH4),  in  which  hydroxyl  is  the  negative  radical. 

The  more  active  bases  are  called  alkalies,  sometimes  caustic 
alkalies  and,  individually,  often,  caustic  potash,  and  caustic  soda. 
The  solutions  are  called  lyes. 

The  name  "base  "  was  originally  applied  to  the  non-volatile,  and  therefore  seem- 
ingly more  fundamental  part  of  a  salt  that  remained  behind  when  the  salt  was 
heated.  Usually  the  negative  radical  is  disintegrated,  as  in  heating  calcium  carbon- 
ate (g.i>.)-  But,  as  a  matter  of  fact,  it  is  generally  the  oxide  and  not  the  hydroxide 
of  the  metal  that  remains.  Still,  the  oxide,  formerly  named  the  base,  often  readily 
gives  the  hydroxide  (cf.  p.  119)  of  which  the  term  "base  "  is  now  used,  and  behaves 
similarly  to  it  in  many  interactions  (cf.  p.  186). 

The  salts  of  hydroxyl  have  the  properties  common  to  all  ionogens. 
They  are  distinguished  from  other  ionogens  by  the  fact  that  their 
negative  radical  is  always  hydroxyl  and  that  in  solution  they  yield 
hydroaridion.  Their  positive  radicals  are  all  different,  K,  Na,  NH4, 
Zo,  Cu,  etc.,  and  constitute  the  positive  ions  of  solutions  of  the  bases. 


THE   CHEMICAL  BEHAVIOR  OF  IONIC   SUBSTANCES         349 

All  the  properties  of  these  ions  in  such  solutions  are  the  same  as 
when  these  ions  are  formed  in  solutions  of  salts  containing  the  same 
radicals.  Usually  the  salts  of  hydroxyl  undergo  double  decomposi- 
tion with  other  ioiiogens,  exchanging  radicals,  even  in  absence  of  a 
solvent.  Like  hydrogen  salts  (p.  347)  they  are  ionized  to  different 
degrees  when  in  solutions  of  equivalent  concentration,  giving  solutions 
with  different  concentrations  of  hydroxidion.  Hence,  they  give 
solutions  of  different  degrees  of  basic  activity. 

The  common  bases,  with  the  exception  of  the  hydroxides  of 
potassium,  sodium,  barium,  strontium,  calcium,  and  ammonium,  are 
but  slightly  soluble  in  water.  Hence,  zinc  hydroxide,  for  example, 
although  it  dissolves  sufficiently  to  enable  chemical  action  to  take 
place  slowly,  does  not  give  enough  hydroxidion  at  one  time  to  affect 
litmus  paper.  Magnesium  hydroxide  and  lead  hydroxide  turn  red 
litmus  paper  blue  with  difficulty.  Doubtless  the  few  molecules  that 
do  dissolve  are  almost  all  ionized : 

Mg(OH)2  (solid)  ±=?  Mg(OH)2  (diss'd)  <=>  Mg"  +  20H', 

but  all  the  dissolved  materials  put  together  (0.01  g.  per  1.)  will 
scarcely  be  weighable  unless  a  considerable  volume  of  the  solution  is 
evaporated. 

Hydroxidion.  —  Hydroxidion  is  a  colorless  substance  found  only 
in  water  and  certain  other  solvents,  and  is  always  associated  with  an 
equivalent  amount  of  some  positive  ion.  It  possesses  a  soapy  taste, 
and  turns  red  litmus  blue  and  colorless  phenolphthalein  pink  (see 
Indicators,  below).  It  confers  great  conducting  power  upon  solutions 
in  which  it  is  contained,  on  account  of  its  speed  of  migratioa  which  is 
second  only  to  that  of  ionic  hydrogen.  It  is  univalent  and  unites  with 
positive  ions. 

Ionic  Double  Decomposition  and  Precipitation  of  Acids  and 
Bases.  —  Under  this  head,  we  shall  discuss  only  the  interaction  of  a 
salt  with  an  acid  or  base,  reserving  for  a  separate  section,  on  neutral- 
ization, that  of  acids  and  bases  with  each  other. 

When  a  highly  dissociated  acid  is  mixed  with  a  salt,  a  reversible 
action  tending  to  form  another  acid  and  salt  is  set  up  (p.  264). 
Such  an  action  is  that  of  nitric  acid  on  a  hypochlorite  (p.  267)  in 
lilute  solution : 

HN08  +  KOC1 1?  KNO,  +  HOC1, 


350  INORGANIC   CHEMISTRY 

giving  potassium  nitrate  and  hypochlorous  acid.  In  such  a  case,  if 
the  products  are  both  as  highly  ionized  as  the  initial  substances,  the 
result  is  similar  to  that  of  the  interaction  between  potassium  chloride 
and  sodium  nitrate  (p.  338).  No  decisive  change  takes  place. 

With  hypochlorous  acid,  however,  which  is  very  slightly  ionized, 
the  result  is  different : 

- 
**11 

This  acid  is  promptly  formed  from  its  ions,  and  the  final  mixture  con- 
tains, mainly,  K*,  NO/  and  molecular  HOC1.  Yet,  since  the  substance 
is  soluble,  no  outward  evidence  that  the  action  differs  from  that  of 
potassium  chloride  and  sodium  nitrate  is  visible.  The  conductivity  of 
the  mixture  (p.  339),  however,  is  found  to  have  been  greatly  reduced 
by  the  removal  of  half  the  ions,  including  the  most  rapidly  migrating 
of  the  four,  hydrion  (p.  314). 

When  the  molecules  of  the  resulting  acid  are  insoluble,  then  it  may 
be  precipitated  (cf.  silicic  acid),  after  the  manner  of  silver  chloride 
(p.  340),  or  may  escape  if  a  gas  (cf.  hydrogen  sulphide),  irrespective 
of  its  degree  of  ionization. 

In  the  same  way,  when  a  salt  and  a  base  are  brought  together,  a 
base  and  a  salt  are  produced.  All  that  has  been  said  in  the  preceding 
paragraph  applies  to  this  case  also.  Thus  ammonium  hydroxide  (q.v.), 
which  is  a  feebly  ionized  base  (p.  331),  is  formed  on  this  plan,  by  mix- 
ing solutions  of  an  ammonium  salt  and  a  strong  base : 

NH4C1  +  NaOH  <=>  NaCl  +  NH4OH. 

When  the  resulting  base  is  insoluble,  like  zinc  hydroxide,  it  is  pre- 
cipitated, and  the  action  becomes  nearly  complete  on  this  account  and 
irrespective  of  the  degree  of  ionization. 

Tonic  Double  Decomposition  and  Activity.  —  It  is  quite  clear 
that  the  complete  formation  of  acids,  bases,  and  salts  by  precipitation 
is  purely  a  result  of  mechanical  details  concerning  solubility,  and 
shows  nothing  about  the  degree  of  affinity  between  the  constituent 
ions.  Again,  the  union  of  ions  to  form  feebly  ionized  substances  only 
shows  the  tendency  of  the  ionic  materials  to  unite  and  may  be  com- 
plete where  the  free  elements  have  little  mutual  activity,  and  vice 
versa.  Thus,  hydrion  and  hypochlorosion  (CIO')  unite  almost  com- 
pletely, while  hydrion  and  chloridion  hardly  unite  at  all.  Yet  hypo- 


THE   CHEMICAL  BEHAVIOR   OF   IONIC   SUBSTANCES         351 

chlorous  acid  is  very  unstable,  while  hydrogen  chloride  is  just  the 
reverse.  Ionic  double  decompositions,  consequently,  give  no  clue  to 
the  activities  of  the  free  materials. 


NEUTRALIZATION. 

Neutralization.  —  When  80  per  cent  sulphuric  acid  is  poured 
upon  solid  potassium  hydroxide,  much  heat  is  developed  and  clouds  of 
steam  arise.  The  solid  product,  when  freed  from  the  rest  of  the  water, 
is  potassium  sulphate.  The  proportions  of  the  materials  used  and  pro- 
duced are  shown  by  the  equation : 

2KOH  +  H2S04 1?  2H20  f  +  K2S04. 

With  any  other  pair  consisting  of  an  acid  and  a  base,  a  similar  interac- 
tion occurs  (cf.  p.  266),  water  and  a  salt  being  produced. 

A  double  decomposition  between  ionogens  is  always  reversible 
(p.  281),  and  so  we  should  expect  that  in  dilute  solution  the  interaction 
of  an  acid  and  a  base  would  be  incomplete.  We  find,  however,  that  this 
particular  sort  of  action  almost  always  goes  so  near  to  completion 
that  it  can  be  employed  for  exact  measurement  of  the  quantity  of  the 
acid  or  base.  This  kind  of  action  is  called  neutralization,  because  both 
acid  and  base  are  completely  consumed,  and  hydrion  and  hydroxidion 
are  alike  impossible  of  detection  in  the  resulting  mixture.  The  solu- 
tion is  neutral  to  litmus. 

Acidimetry  and  Alkalimetry.  —  If  the  problem  is  to  ascertain 
the  weight  of  hydrogen  chloride  in  each  liter  of  a  specimen  of  hydro- 
chloric acid,  this  can  be  done  by  neutralizing  a  measured  portion  of 
this  acid  with  a  solution  of  an  alkali  of  known  concentration.  The 
volume  of  the  latter  which  is  required  for  the  purpose  is  observed.  If 
the  alkali  is  sodium  hydroxide,  the  action  taking  place  is  : 

HC1  +  NaOH  ->  H20  +  NaCl. 

The  volume  of  acid  is  measured  out  into  a  beaker  by  means  of  a 
pipette  (Fig.  82)  of  fixed  capacity,  which  is  filled  by  suction  to  the 
mark  on  the  stem.  Suppose  the  amount  to  be  25  c.c.  The  standard 
(cf.  p.  236)  alkali  solution  is  placed  in  a  burette  (Fig.  83),  which  is 
filled  down  to  the  tip  of  the  nozzle.  A  few  drops  of  litmus  solution 
are  now  added  to  the  acid,  and  the  alkali  is  allowed  to  run  in  slowly. 
After  a  time,  the  hydroxidion  which  this  introduces  will  begin  to  pro- 


352 


INORGANIC   CHEMISTRY 


duce  a  blue  color  close  to  where  the  stream  enters  the  liquid.     This  is 
at  first  dissipated  by  stirring,  and  the  whole  remains  red.     Finally, 

however,  a  point  is  reached  at  which  the  entire  solution  assumes 

a  tint  intermediate  between  blue  and  red.     With  one  drop  less 

of  the  base,  it  is  distinctly   red.     With  one   drop   more,  it 

would  become  distinctly  blue.     Litmus  paper  of  either  shade 

dipped  in  this   neutral   solution  remains   unaffected.     It   is 

needless  to  say  that  the  acid  might  have  been  added  to  the 

base  with  the  same  final  result. 

The  standard   solutions   used   in   this   work  are  usually 

normal,  and  contain  one  equivalent  weight  of  the  alkali  or 

acid  in  one   liter  of  the  solution.     For  more  delicate  work, 

deci-normal     solutions 

may  be  employed.    The 

concentration  of    such 

a  solution  is  called  its 

tiver,  and  the  operation 

of  neutralizing  another 

solution    by  means  of 

it,  titration.  The  value 

of  standard    solutions 

lies   in   the  fact  that 

when  once  the  solution 

has  been  prepared,  and 

the  exact  concentration 

adjusted  by   quantita- 
tive   experiments,    its 
PIG  82.   use   does    not   require 

any  weighing,  and  the 
measurements  of  volumes  can 
be  carried  out  with  great  rapid- 
ity. A  process  involving  weigh- 
ing need  not  again  be  under- 
taken until  the  stock  of  the 
standard  solution  is  exhausted. 
The  calculation  of  the  result 
is  also  simple.  One  liter  of  FlG  ^ 

normal  alkali  contains  17  g.  of 

available  hydroxyl,  and  one  liter  of  normal  acid  1  g.  of   available 
hydrogen.     Equal  volumes  of  normal  solutions  will  therefore  exactly 


THE   CHEMICAL  BEHAVIOR   OF  IONIC   SUBSTANCES         353 

neutralize  one  another,  18  g.  of  water  being  formed  by  interaction  of 
a  liter  of  each.  If,  for  the  neutralization  of  the  25  c.c.  of  hydrochlo- 
ric acid  used  above,  50  c.c.  of  normal  alkali  are  required,  the  acid  is 
twice-normal  (2N).  When  15  c.c.  are  required,  the  acid  is  £f  or  £  N. 
If  the  actual  weight  of  hydrogen  chloride  in  the  latter  case  has  to  be 
calculated,  we  remember  that  there  are  36.45  g.  of  the  compound  in 
1  1.  of  a  normal  solution,  and  therefore  36.45  X  f  X  T§£^  g.  =  .5467  g. 
in  25  c.c.  of  one  which  is  £  normal. 

Methods  of  quantitative  analysis  in  which  standard  solutions  (cf. 
pp.  148,  236,  307)  are  employed  are  known  as  volumetric  methods,  and 
are  much  used  by  analysts  and  investigators.  They  occupy  much  less 
time  than  gravimetric  operations,  in  which  numerous  weighings  have 
to  be  made,  and  are  often  just  as  accurate.  The  substances,  like 
litmus,  by  whose  change  of  color  the  completeness  of  the  action  is 
made  known,  are  called  indicators  (see  below). 

Theory  of  Neutralization.  —  The  neutral  mixture  of  the  acid 
and  base  gives  no  evidence  of  the  presence  either  of  the  hydrogen  ions 
or  of  the  hydroxyl  ions.  The  characteristic  tastes,  and  actions  upon 
indicators,  of  these  two  ions,  and  the  interaction  of  the  former  of  the 
two  with  metals  like  magnesium,  are  all  wanting.  That  this  is  due, 
not  simply  to  two  opposing  influences  having  destroyed  each  other's 
effects,  but  to  a  real  disappearance  of  the  agencies  themselves,  may  be 
demonstrated  by  showing  that  the  total  number  of  ions  is  very  much 
smaller  in  the  mixture  than  in  the  two  substances  taken  separately. 
The  trough  (Fig.  81,  p.  327)  is  half-filled  with  a  dilute  solution  (say,  N/4) 
of  some  active  acid,  such  as  hydrochloric  acid.  An  equal  volume  of 
a  N/4  solution  of  some  soluble  base,  such  as  sodium  hydroxide,  is 
then  allowed  to  flow  in,  below  the  acid.  On  completing  the  circuit  we 
find  a  considerable  deflection  of  the  amperemeter  (say,  1.5  amperes). 
When  the  interaction  is  now  brought  about  by  stirring,  a  very  great 
fall  in  the  reading  (say  to  0.5  amperes)  is  observed.*  The  only 
plausible  explanation  is  that,  not  only  have  many  of  the  ions  assumed 
a  molecular  form,  but  those  which  have  suffered  in  this  respect  have 
been  the  most  rapidly  moving  and  best  conducting  ones,  namely,  the 
hydrion  and  hydroxidion. 

*  The  experiment  may  be  made  more  striking  by  adding  a  few  drops  of  phenol- 
phthalein  solution  to  the  acid  and  using  a  minute  excess  of  the  base.  To  prevent 
the  appearance  of  a  pink  layer  at  the  interface,  and  before  the  stirring,  a  thin  layer 
of  sodium  chloride  solution  may  be  introduced  below  the  acid,  before  the  layer  of 
the  base  is  added. 


354  INORGANIC   CHKMISTAY 

The  general  plan  of  all  interactions  of  acids  and  bases  is  as 
follows : 

HC1 1?  Cl'  +  H- 
NaOHi=>Na'+OH' 

it        it 
NaCl  H20 

The  ionization  of  the  hydrochloric  acid  reaches  0.785  in  a  normal 
solution,  and  goes  further  when  the  acid  is  diluted  with  the  water  of 
another  solution.  That  of  the  sodium  hydroxide  similarly  goes  beyond 
0.73.  Thus  the  initial  substances  are  almost  entirely  ionic.  The 
crosswise  union,  H*  -f  OH'  t=;  H2O,  however,  is  all  but  complete,  for 
water  is  hardly  ionized  at  all  (p.  331).  The  materials  on  whose  inter- 
action with  the  Cl'  and  Na*,  respectively,  the  maintenance  of  mole- 
cules HC1  and  NaOH  depends,  being  thus  removed,  the  dissociation  of 
the  acid  and  base  promptly  brings  itself  to  completion,  and  the  left 
sides  of  the  equations  vanish.  Practically  all  the  hydrion  and 
hydroxidion  become  water.  The  Cl'  and  Na',  however,  if  the  solution 
is  now  semi-normal,  unite  to  the  extent  of  0.266  only  (p.  331).  If 
it  is  more  dilute,  this  union  forms  a  still  smaller  factor  in  the  whole 
change.  Practically  it  is  negligible.  Now  all  that  has  been  said  of 
this  acid  and  base  will  apply  mutatis  mutandis  whenever  any  active, 
highly  ionized  acid  and  base  come  together.  Thus  we  may  write 
one  simple  equation  for  all  neutralizations  of  active  acids  and  bases  : 

H'  +  OH'  -»  H2O, 

without  omitting  anything  essential. 

The  ions  of  a  salt  are  always  left  over  from  the  main  action,  and 
may  be  brought  together,  in  turn,  by  evaporation  (cf.  p.  341). 

The  equations  as  commonly  written  : 

NaOH  +  HC1  — >  NaCl  +  H2O, 
2NaOH  +  H2S04  ->  Na2SO4  +  2H2O, 

apply  to  the  interactions  when  water  is  absent.  If  used  for  neutralization  in 
dilute  solution,  it  must  be  understood  that  they  condense  two  changes  into  one 
equation.  The  formation  of  water  comes  first,  that  of  the  salt  afterwards.  Some- 
times neutralization  is  wholly  misconstrued  by  the  supposition  being  made  that  it 
occurs  in  consequence  of  a  great  tendency  to  salt  formation. 

It  will  be  seen  that  neutralization  is  the  precise  reverse  of  hydroly- 
sis (pp.  181,  344).  The  former  being  almost  always  nearly  complete, 
the  latter  must  be,  as  a  rule,  very  slight 


THE   CHEMICAL  BEHAVIOR   OF   IONIC   SUBSTANCES         355 

Indicators.  —  Indicators  are  substances  which,  in  presence  of 
certain  other  substances,  assume  a  very  deep  color,  or  change  sharply 
from  one  deep  color  to  another.  Thus,  phenolphthale'm  (p.  349)  is 
colorless  in  presence  of  acids  (i.e.,  hydrion),  and  red  (when  dilute 
pink)  in  presence  of  alkalies  (i.e.,  hydroxidion).  Litmus,  again,  is  red 
with  acids,  and  blue  with  alkalies.  The  change  of  color  depends  upon 
a  chemical  interaction  in  each  case,  but  since  indicators  are  chosen  for 
their  strong  coloration,  the  quantity  of  the  acid  or  base  used  up  in  chang- 
ing the  tint  of  the  trace  of  the  indicator  is  so  small  as  to  be  negligible. 

The  common  indicators  are  : 

Phenolphthalein,  C14H1004,  a  colorless  substance  and  very  feeble 
acid.  It  is  not  perceptibly  dissociated  into  its  ions : 

C14H1004  (colorless)  ±?  C14H90/  (red)  +  H*, 

and  in  neutral  or  acid  solutions  is,  therefore,  without  visible  color. 
When  a  base  is  added  gradually  to  an  acid  containing  some  of  this 
indicator,  the  acid  is  first  neutralized.  Then,  and  not  till  then,  the 
slightest  excess  of  hydroxidion  unites  with  the  trace  of  hydrion  from 
the  phenolphthalein,  the  above  equilibrium  is  displaced  forwards,  and 
a  visible  amount  of  the  red  negative  ion  is  formed : 

C14H1004  (colorless)  «=;  C14H,0/  (red)*  +     H*) 
NaOH  fcsNa'  +OH'$- 

This  indicator  shows  the  presence  of  an  excess  of  alkali  most  sharply 
when  the  alkali  is  an  active  one  like  sodium  hydroxide,  and  should, 
therefore,  be  employed  only  with  strong  bases.  With  a  weak  base 
like  ammqnium  hydroxide,  a  considerable  excess  must  often  be  used 
before  the  color  appears. 

Litmus  is  a  natural  dyestuff  of  unknown  chemical  structure. 
Doubtless,  however,  one  of  its  colors  is  that  of  the  molecule,  and  the 
other  that  of  the  ion. 

Methyl  orange,  (CH8)2NC6H4.N  :  KC6H4SOg]Sra,  is  a  complex  organic 
compound  which  gives,  in  acid  solution,  a  red  and  in  alkaline  solution 
a  yellow  color. 

Congo  red  is  the  sodium  salt  of  an  acid  of  complex  structure  (see 
Dyes).  In  neutral  or  alkaline  solutions  it  is  red;  with  acids  it  turns 

*  The  ion  has  this  composition,  but,  in  reality,  has  a  different  chemical  struc- 
ture from  the  corresponding  part  of  the  original  molecule.  An  internal  rearrange- 
ment, not  representable  in  the  equation,  accompanies  the  dissociation.  The  same 
remark  applies  to  the  other  indicators. 


356  INORGANIC   CHEMISTRY 

blue.  Paper  dipped  in  Congo  red  differs  from  litmus  paper  in  that  it 
shows  gradations  in  color,  the  blue  being  much  more  distinct  with  an 
active  acid  than  with  a  relatively  weak  one  like  acetic  acid  (p.  347). 
Litmus  paper  is  equally  red  with  all  acids  save  the  very  feeblest. 

Some  special  indicators  have  been  mentioned.  Thus,  starch  emul- 
sion is  used  for  recognizing  the  presence  of  traces  of  iodine  (p.  235). 
Potassium  permanganate  is  itself  so  strongly  colored  that  it  is  its  own 
indicator  (p.  307). 

Neutralization  of  Little  Ionized  Substances.  —  When  concen- 
trated solutions  are  employed,  or  acids  and  bases  which  are  but  little 
ionized  are  involved,  the  mechanism  of  the  change  is  still  the  same  in 
all  respects.  The  only  difference  is  that,  since  the  acid  or  base  is  not 
fully  ionized  to  start  with,  its  molecules  must  dissociate  progressively, 
in  proportion  as  the  hydrogen  ions  pass  into  combination.  All  the 
hydrogen  and  hydroxyl  capable  of  forming  ions  will  pass  through  that 
stage  and  ultimately  become  water  before  the  solution  can  reach  the 
neutral  condition. 

From  this  it  will  be  seen  that  the  activity  of  acids  and  bases  cannot 
be  measured  by  the  quantity  of  base  or  acid  required  to  neutralize 
them.  The  full  amount  required  by  the  equation  is  always  needed 
in  every  case.  This  is  because  neutralization  uses  up  the  hydrion  or 
hydroxidion  at  once,  and  so  permits  the  rapid  generation  of  a  fresh 
supply.  The  concentration  of  one  of  these  ionic  materials  can  only  be 
measured  by  some  action  which  uses  it  up  slowly  or  not  at  all,  so  that 
ionic  double  decompositions  are  excluded.  In  the  action  of  metals  on 
acids  (p.  347)  and  in  determining  conductivity  (p.  325)  the  consumption 
of  the  ions  is  slow,  and  hence  the  measurement  can  be  made  in  these 
cases.  Actions  which  consume  no  ions  at  all  are  also  known,  and  are 
used  in  measuring  activity  (see  Carbohydrates  and  esters). 

When  the  acid  or  base  is  but  little  soluble  in  water,  as  when  zinc 
hydroxide  is  treated  with  a  dilute  acid,  one  other  link  is  added  to  the 
network  of  equilibria.  The  acid  proceeds  to  interact  with  the  small 
dissolved  part  of  the  base.  As  this  is  disposed  of,  solution  goes  on 
progressively,  and,  through  a  train  of  equilibria : 

Zn(OH)2  (solid)  t=»  Zn(OH)2  (diss'd)  +±  Zn"+  20H', 

the  supply  of  hydroxidion  is  maintained  until  all  the  molecules  of  the 
base,  solid  and  dissolved,  are  used  up  and  the  action  is  completed. 
Heating  hastens  these,  as  it  does  all  other  changes. 


THE   CHEMICAL  BEHAVIOR   OF   IONIC   SUBSTANCES         357 

If  acid  and  base  are  alike  insoluble,  it  is  best,  if  the  production  of 
the  salt  is  the  ultimate  object,  to  fuse  the  materials  together  at  a  high 
temperature. 

Thermochemistry  of  Neutralization.  —  The  above  interpreta- 
tion of  the  phenomena  of  neutralization  is  confirmed  by  many  facts. 
Thus  a  considerable  amount  of  heat  is  liberated  in  neutralization. 
Now,  when  active  acids  (p.  347)  and  bases  (p.  349)  in  dilute  solution 
are  concerned,  it  is  found  that  the  quantities  of  heat  for  the  neutraliza- 
tion of  the  same  amount  of  hydrion,  or  hydroxidion,  are  always  the 
same,  namely,  13,700  cal.  for  equivalent  weights.  If  the  action  con- 
sisted primarily  in  the  formation  of  a  different  salt  from  every  pair, 
we  should  expect  the  heat  liberated  to  be  different.  Thus,  the  heats  of 
formation  of  dry  potassium  chloride  and  dry  sodium  iodide  are  10,120 
cal.  and  7030  cal.,  respectively.  But  the  heats  of  formation  of  their 
solutions  by  neutralizing  the  proper  acids  and  bases  are  identical.  If, 
however,  in  such  cases,  neutralization  consists  always  simply  in  the 
formation  of  water,  we  should  expect  the  quantities  of  heat  liberated  to 
be  identical,  as,  in  fact,  they  are : 

H"  +  OH'  --»  H20  +  13,700  cal. 

We  are  confirmed  in  these  conclusions  when  we  employ  concentrated  solutions, 
or  use  less  completely  ionized,  or  insoluble  acids  and  bases  for  neutralization.  With 
such  substances, — and  they  are  in  the  majority, — the  heats  of  neutralization  are  not 
alike,  but  different  in  every  case.  Thus,  for  dilute  solutions  of  sodium  hydroxide 
and  hydrofluoric  acid,  the  latter  a  slightly  ionized  soluble  acid,  the  thermochemical 
equation  is  as  follows  : 

NaOH  +  HP  -*  H2O  +  NaF  +  16,270  cal. 

Since  the  sodium  fluoride  is  fully  ionized,  the  only  difference  between  this  case  and 
the  preceding  one  is  that  the  hydrogen  fluoride  is  largely  in  the  molecular  condi- 
tion to  start  with,  and  that  here,  in  addition  to  the  union  of  hydrogen  and 
hydroxyl  ions,  we  have  a  continuous  dissociation  of  the  hydrofluoric  acid  accom- 
panying the  neutralization.  The  fact  that  here  the  heat  produced  is  much  greater 
than  before,  shows  that  the  dissociation  of  this  acid  is  associated  with  the  produc- 
tion of  heat  (cf.  pp.  260,  330).  When  the  same  base  is  used  with  hypochlorous  acid, 
the  divergence  is  in  precisely  the  opposite  direction  and  about  the  same  in  amount: 

NaOH-f-HCIO  ->  H2O  +  NaCIO  +  9840  cal. 

Here  again  the  salt  produced,  sodium  hypochlorite,  is  fully  ionized,  so  that  the  dim- 
inished evolution  of  heat  must  be  due  to  the  fact  that  the  feebly  ionized  hypochlo- 
rous acid  absorbs  part  of  the  heat  of  neutralization  in  passing  into  the  ionic 
condition.  Applying  this  to  bases,  we  find  that  the  neutralization  of  ammonium 
hydroxide,  a  feebly  ionized  base,  with  any  active  acid,  gives  a  heat  of  neutraliza- 
tion below  the  normal : 


358  INORGANIC   CHEMISTRY 

NH4OH  +  HC1  ->  H.,0  +  NH4C1  +  12,200  cal. 

Here  again  the  salt  produced  is  fully  ionized.  Thus  the  ionization  of  the  ammo- 
nium hydroxide  must  have  consumed  an  appreciable  part  of  the  heat  of  neutraliza- 
tion which  would  otherwise  have  reached  the  normal  figure  of  13.700  calories. 

Volume  Change  in  Neutralization.  — When  the  volumes  of  the  solu- 
tions of  active  acids  and  alkalies  are  carefully  measured  before  being  mixed,  and 
compared  with  the  volume  of  the  neutral  mixture,  an  expansion  is  always  found  to 
have  occurred.  When  one  liter  of  a  normal  solution  of  each  substance  is  taken  at 
starting,  the  volume  of  the  mixture  is  always  20  c.c.  greater  than  that  of  the  com- 
ponent liquids.  When  less  highly  ionized  acids  and  bases  are  used,  the  alteration 
in  volume  is  irregular,  since  it  is  affected  by  the  occurrence  of  other  changes  than 
the  mere  union  of  hydrion  and  hydroxidion. 

MIXED    lONOGENS    AND    DOUBLE    SALTS. 

As  a  rule,  a  univalent  ion,  such  as  chloranion  (CIO/),  unites  with 
one  kind  of  cation  and  gives  but  one  kind  of  salt  (c/.,  however,  p.  242). 
The  result  is  called  a  neutral  or  normal  salt,  as  KC1O,  or  NaC108. 
The  acid,  chloric  acid,  is  called  a  monobasic  acid,  for  its  molecule 
reacts  with  but  one  molecule  of  a  base.  The  possibilities  are  more 
numerous,  however,  with  an  ion  of  higher  valence.  Thus : 

CABBONANION  MAY  GlVK  t  CALCION  MAY  GlVB  : 

H2C08,  the  acid,  Ca(OH)2,  the  base, 

NajjC08,  a  neutral  salt,  CaCl2,  a  neutral  salt, 

NaHC08,  an  acid  salt,  Ca(OH)Cl,  a  basic  salt,* 

NaKC08,  a  mixed  salt,  CaCl(OCl),  a  mixed  salt. 

Carbonic  acid  is  a  di-basic  acid,  and  calcium  hydroxide  a  di-acid  base. 
The  last  two  compounds  of  each  set  are  mixed  ionogens.  Their  char- 
acteristic is  that  they  contain  more  than  two  kinds  of  radicals  and 
break  up  in  solution,  giving  more  than  two  kinds  of  ions. 

Acid  Salts.  —  The  acid  salts  are  obtained  by  using  half  that  quan- 
tity of  the  base  which  would  be  required  fully  to  neutralize  the  acid,  and 
evaporating  the  resulting  solution  : 

NaOH  +H2S04  *=?  H2O  |  +  NaHS04. 

With  a  monobasic  acid,  say  hydrochloric  acid,  this  treatment  gives 
simply  a  mixture  of  the  normal  salt  and  the  free  acid,  and  not  a  single 
substance. 

»  This  particular  basic  salt  has  not  been  isolated  in  a  pure  state. 


THE   CHEMICAL  BEHAVIOR  OF  IONIC   SUBSTANCES         359 

Acid  salts  are  also  formed  by  the  interaction  of  other  salts  with  an 
excess  of  the  acid  (pp.  178,  179,  227). 

The  acid  salt  is  intermediate  in  composition  between  the  acid  itself 
and  the  normal  salt.  All  of  the  hydrogen  of  the  acid  has  not  been 
displaced  by  the  metal.  It  is  named  an  acid  salt  on  account  of  its 
composition,  but  is  not  necessarily  acid  in  its  reaction  towards 
litmus.  That  depends  on  whether  its  solution  contains  a  sufficient 
amount  of  hydrion  to  affect  indicators;.  Sodium  hydrogen  sulphate 
gives  the  ions  Na*  and  HSO/,  but,  even  in  moderately  dilute  solution, 
the  latter  ion  is  further  dissociated  into  H*  and  SO/'  to  a  large 
extent  (p.  346).  Its  solution  is  therefore  acid  in  truth.  On  the  other 
hand,  sodium  hydrogen  carbonate,  NaHC03,  derived  from  carbonic  acid, 
H2C08,  gives  the  ions  Na  *  and  HCO/,  and  the  amount  of  hydrion 
formed  by  the  latter  is  too  small  to  be  detected  by  indicators.  This 
acid  salt  gives  therefore  a  solution  which  is  actually  neutral  to  litmus.* 

Basic  Salts.  —  Corresponding  to  the  acid  salts  we  have  also  basic 
salts,  about  which  statements  parallel  to  the  above  might  be  made. 
Thus,  from  sodium  hydroxide  but  one  salt  can  be  formed.  With  lead 
hydroxide,  Pb(OH)2,  however,  the  displacement  of  one  hydroxyl  by  a 
negative  radical,  without  the  disturbance  of  the  other,  is  conceivable 
and  can  be  achieved.  The  half -chloride,  for  example,  is  called  lead 
oxychloride  (Pb(OH)Cl).  The  basic  salts  are  usually  insoluble  in 
water,  and  therefore  as  a  rule  do  not  exhibit  the  basic  reaction  with 
litmus. 

Mixed  Salts.  —  So-called  mixed  salts,  like  sodium-potassium  car- 
bonate KNaC08,  (see  Silicates),  may  be  obtained  by  half  neutralizing 
the  acid  with  the  calculated  amount  of  one  base  and  then  completing 
the  operation  with  the  other.  Corresponding  treatment  will  give 
mixed  salts  of  a  di-acid  base. 

It  will  be  seen  that  2KNaC08  is  equivalent  to  K2C08,]Sra2C08,  and 
that  2CaCl(OCl)  is  equivalent  to  CaCl2,Ca(OCl)2.  Since  we  have  as 
yet  no  general  means  of  determining  the  molecular  weights  of  solids, 
there  is  no  generally  applicable  way  of  deciding  which  formula  is 
preferable  (see,  however,  Bleaching  powder,  p.  266,  and  under  Calcium). 
In  solutions  of  these  salts  the  ions  which  are  found  might  come  from  a 

*  Because  of  hydrolysis  (p.  344),  the  solution  of  the  "neutral"  salt,  sodium 
carbonate  Na2CO3,  is  actually  alkaline  in  reaction.  The  terms  "  acid,"  "basic,"  and 
"  neutral,"  applied  to  salts,  refer  simply  to  the  composition  and  ignore  the  behavior. 


360  INORGANIC   CHEMISTRY 

substance  possessing  either  of  the  alternative  formulae,  so  that  no  light 
is  thrown  on  the  question  by  this  means.  Thus,  most  compounds  of  this 
kind,  with  the  exception  of  acid  and  basic  salts,*  are  considered  to  be 
molecular  compounds  (p.  123)  of  two  salts  and  are  classed  as  double  salts. 

Double  Salts.  —  Substances  similar  to  ferrous-ammonium  sulphate 
FeS04,(NH4)2S04,6H20,  are  very  numerous.  Because  their  formulae 
can  be  written  so  as  to  show  two  complete  salts,  and  because  they  are 
easily  formed  by  crystallization  from  a  solution  containing  both  salts, 
they  are  called  double  salts.  In  solution  they  are  resolved  into  their 
constituent  salts,  and  these,  in  turn,  are  ionized.  Almost  always  the 
acid  radicals  are  identical  (see,  however,  Kainite). 

Each  kind  of  ion  of  a  double  salt  exhibits  its  own  properties,  irre- 
spective of  the  nature  of  the  numerous  substances,  ionic  and  otherwise, 
which  are  present.  Hence,  when  a  solution  of  a  particular  ionic 
material  is  required,  solutions  of  such  bodies  are  often  used  instead  of 
those  of  simpler  ones,  if  for  any  reason  the  substitution  is  convenient. 
The  choice  of  the  complex  compound  must  be  made  in  such  a  way 
that  the  other  ions  shall  not  interfere  with  that  particular  reaction  of 
one  of  them  which  is  in  question. 

The  class  of  bodies  known  as  salts  of  complex  acids  (q.v.)  are 
ionized  like  ordinary  salts  and  not  like  double  salts. 

KINDS  OF  IONIC  CHEMICAL  CHANGE. 

Five  distinct  varieties  of  chemical  change  are  characteristic 
of  ionic  materials.  These  are  :  (1)  Disunion  or  combination  of  ions, 
(2)  displacement  of  the  material  of  one  ion  by  another  substance,  (3) 
destruction  or  formation  of  a  compound  ionic  material,  (4)  change  in 
the  charges  of  two  ionic  materials,  (5)  charge  or  discharge  of  two  ionic 
materials,  electrically.  Every  one  of  these  kinds  of  action  has  been 
illustrated,  some  of  them  very  frequently,  in  the  present  and  foregoing 
chapters. 

Disunion  and  Combination  of  Ions.  —  This  sort  of  change  is 
illustrated  every  time  an  ionogen  is  dissolved  in  water  (disunion)  or  a 
solution  of  such  a  substance  is  evaporated  (combination).  Both  of  the 

*  The  formulae  of  basic  salts  even  are  often  written  as  if  they  were  molecular 
compounds,  as  Cu(OH)2,CuCl2,  or  even  CuO,CuCl2,H2O,  in  place  of  Cu(OH)C) 
(see  Copper). 


THE   CHEMICAL   BEHAVIOR   OF   IONIC   SUBSTANCES         3(51 

directions  of  this  sort  of  change  occur  also  to  a  greater  or  less  extent 
whenever  solutions  of  two  ionogens  are  mixed.  In  the  latter  case  : 

(1)  Two  salts  give  two  salts  (pp.  337-341). 

(2)  An  acid  and  salt  give  a  salt  and  an  acid  (p.  349). 

(3)  A  base  and  salt  give  a  salt  and  a  base  (p.  350). 

(4)  An  acid  and  base  give  water  and  a  salt  (neutralization). 

(1)  is  complete  only  when  at  least  one. product  is  insoluble.  (2) 
and  (3)  are  complete  when  at  least  one  product  is  little  ionized  or 
insoluble  or  both.  (4)  is  almost  always  complete  because  water  is 
generally  less  ionized  than  any  other  substance  in  the  system. 

Displacement  of  One  Ion  :  Electromotive  Series  of  the  Metals. 

-  When  zinc,  aluminium,  and  other  metals  are  placed  in  a  dilute  acid, 
hydrogen  is  liberated  (pp.  95  and  346) : 

Zn  +  2H--->Zn"-f-H2. 

This  action  takes  place  with  all  acids,  because  it  concerns  in  reality 
only  the  hydrion  in  the  solution.  Its  speed  for  any  one  metal  depends 
on  the  concentration  of  hydrion  (p.  347).  Similarly,  fluorine  displaces 
chlorine  from  chloridion  (p.  342),  chlorine  displaces  bromine  from 
bromidion  (p.  228)  : 

Cl2+2Br'-»2Cl'+Br2, 

and  bromine  displaces  iodine  from  iodidion  (p.  236).  Each  of  these 
actions  is  independent  of  the  nature  of  the  other  ion  which  accompa- 
nies the  one  undergoing  change. 

The  same  sort  of  displacement  occurs  with  all  positive  ions.  Thus, 
zinc  will  not  only  displace  hydrogen,  but  also  other  metallic  elements, 
like  iron,  lead,  copper,  and  silver,  from  the  ionic  condition  in  solutions  of 
their  salts.  Lead,  in  turn,  will  displace  copper  and  silver,  but  not  zinc  or 
iron.  Copper  will  displace  silver.  Thus  the  metals  can  be  set  down 
in  an  order  such  that  each  metal  displaces  those  following  it  in  the 
list  and  is  displaced  by  those  preceding  it.  This  list  (see  next  page) 
is  known  as  the  electromotive  series  of  the  metals,  because  in  electro- 
lysis of  normal  solutions  of  their  salts,  the  electromotive  force  of 
the  current  required  to  deposit  each  metal  (cf.  p.  324)  is  less  than 
that  for  the  metal  preceding  in  the  list  (see  Electromotive  chemis- 

try). 

This  list  embodies  many  facts  in  the  behavior  of  the  metals,  and 


362 


INORGANIC   CHEMISTRY 


should  be  kept  in  mind  as  furnishing  a  key  to  the  actions  in  which  a 
free  metal  is  used  or  produced.  For  example,  the  chemical  activity 
of  the  free  metals  places  them  in  the  same  order.  The  earliest  ones 
rust  much  more  readily  in  air  than  do  the  later  ones.  Those  following 
copper  do  not  rust.  Conversely,  the  oxides  of  the  metals  down  to 
and  including  manganese,  when  heated  in  a 
stream  of  hydrogen,  may  give  lower  oxides,  but 
are  not  completely  reduced.  The  oxides  of  cad- 
mium and  succeeding  metals  are  easily  reduced. 
The  oxides  of  mercury  and  the  last  four  metals 
are  decomposed  by  heating  alone.  The  relations 
of  the  metals  in  respect  to  combination  with 
elements  other  than  oxygen  are  similarly  ex- 
pressed by  the  arrangement  in  this  table. 

The  position  of  hydrogen  is  particularly  sig- 
nificant. It  will  be  noted  that  none  of  the  metals 
preceding  hydrogen  are  found  free  in  nature  as 
ordinary  minerals,*  while  all  of  the  metals  suc- 
ceeding hydrogen,  although  occurring  to  some 
extent  in  combination,  are  found  also  free.  The 
explanation  of  this  is  that,  by  prolonged  action 
upon  ordinary  water,  containing,  as  it  must,  car- 
bonic acid  and  other  sources  of  hydrogen  ions, 
the  metals  preceding  hydrogen  must  eventually 
displace  hydrion  and  pass  into  some  form  of 
combination  (cf.  p.  346).  The  metals  following 
hydrogen  do  not  displace  hydrion  and  are  much 
less  affected  by  the  agencies  which  are  most 
active  in  the  chemical  transformation  of  minerals. 
Hence  they  often  remain  in  the  free  state. 
The  negative  Ions  can  be  arranged  in  order  in  a  similar  way.f 


ELECTROMOTIVE 

SERIES  OF  THE 

METALS. 

Alkali  metals  (g.v.) 
Alkaline  earth 
metals  (q.v.) 
Magnesium 
Aluminium 
Manganese 
Zinc 

Chromium 
Cadmium 
Iron 
Cobalt 
Nickel 
Tin 
Lead 

Hydrogen 
Copper 
Arsenic 
Bismuth 
Antimony 
Mercury 
Silver 
Palladium 
Platinum 
Gold 


*  Free  lead  and  tin  do  occur  as  rare  minerals.  Iron,  with  a  little  cobalt  and 
nickel,  constitutes  many  meteoric  masses. 

t  To  avoid  a  common  misconception,  it  must  be  noted  that  the  electromotive 
series  has  no  bearing  on  the  tendency  of  one  radical  to  dislodge  another  in  double 
decompositions.  The  place  of  an  element  in  the  E.M.  series  defines  its  relative 
activity  when  free,  and  has  to  do  only  with  actions  where  one  free  element  dis- 
places (p.  99)  another.  The  influences  which  determine  a  double  decomposition 
(cf.  pp.  180,  257-260)  have  little  to  do  with  the  chemical  activity  of  the  com- 
pounds concerned  (p.  350),  and  nothing  at  all  to  do  with  that  of  the  free  elements, 
for  these,  in  fact,  are  not  present  at  all  (see  Halides  of  silver). 


THE  CHEMICAL  BEHAVIOR  OF  IONIC   SUBSTANCES         363 

Destruction  or  Formation  of  a  Compound  Ion.  —  The  de- 
struction of  a  compound  ionic  material  is  observed  in  the  action  of  any 
reducing  agent,  such  as  hydrogen  peroxide  (p.  307),  upon  a  dilute 
solution  of  a  permanganate.  The  compound  ion  MnO/  gives  by 
reduction  Mn"  and  water.  It  was  also  encountered  (p.  278)  in  the 
reduction  of  I08'  by  hydriodic  acid,  in  which  free  iodine  and  water 
were  formed: 

H-  +  10,'  +  5H'  +  51'  ->  3H20  +  3I2. 

The  converse  occurs  when  potassium  cyanide  is  added  in  excess  to 
a  solution  of  a  salt  of  silver.  First,  silver  cyanide  is  precipitated,  and 
then  this  compound  unites  with  the  excess  of  cyanidion : 

Ag-  +  NO.'  +  K-  +  NO'  ^  K'  4-  N0'8  +  AgNC  J, 
K-  4-  ON' 4-  AgCNf->  K-  +  Ag(CN)/. 

The  product  is  the  soluble  potassium  argent  icy  anide.  It  is  a  salt  of 
the  complex  acid  HAg(CN)2,  and  not  a  double  salt  (p.  360).  It  does 
not  decompose  into  potassium  and  silver  cyanides  and  their  ions  when 
in  solution,  for  the  second  action,  above,  is  not  appreciably  reversible. 

Change  in  the  Charges  of  Two  Ions.  —  A  reduction  in  the 
charge  in  two  ions  probably  occurs  in  the  preparation  of  chlorine 
(p.  171).  The  decomposition  of  the  manganese  tetrachloride,  if  it  is 
indeed  formed,  takes  place  by  the  simultaneous  discharge  of  two 
equivalents  of  electricity  from  the  quadrivalent  manganese  ion  and 
two  ions  of  chloridion : 

Mn'—  -f  401'  -»  Mn-  +  2C1'  +  C12. 

Both  this  sort  of  change  and  its  converse  are  common  with  ions  of 
metals  such  as  iron  and  tin  (g.v.),  which  have  more  than  one  valence. 

Charge  and  Discharge  of  Tivo  Ions,  Electrically.  —  Dis- 
charge of  ions  is  brought  about  in  every  electrolysis  (p.  318).  Thus, 
when  hydrochloric  acid  is  decomposed  by  the  current,  we  have  : 

2H-f20->H2      and      2C1  +  2  ® ->  CL, . 

The  converse  takes  place  when  the  polarization  current  (p.  324)  is 
allowed  to  flow.  Both  charge  and  discharge  occur  in  every  simple 
battery,  as  when  zinc  dissolves  in  dilute  sulphuric  acid  to  give  zinc 
sulphate  (p.  346) : 

Zn-»Zn  +  2t-)      and       2H->H   +  20. 


364  INORGANIC   CHEMISTRY 

The  creation  of  the  positive  charges  in  the  former  of  the  two 
equations  leaves  the  rod  of  zinc  negatively  charged,  that  of  the  nega- 
tive charges  in  the  latter  renders  the  platinum  wire  positive. 

Two  or  More  Kinds  of  Ionic  Action  Simultaneously.  —  In 

the  interaction  of  iodic  acid  and  hydrogen  iodide,  both  the  decomposi- 
tion of  a  compound  ion  (IO/)  and  the  change  in  charge  on  other  ions 
(I'  giving  I2,  and  H*  giving  water)  occur  simultaneously.  Compli- 
cations of  this  kind  are  not  uncommon. 

NON-IONIC  MODES  OF  FORMING  IONOGENS. 

While  ionogens  may  always  be  made  by  the  union  of  the  proper 
ions,  they  must  nevertheless,  in  the  absence  of  the  solvent,  be  regarded 
as  chemical  substances  which  may  he  constructed  out  of  their  constit- 
uents without  reference  to  the  ionic  plane  of  cleavage.  Thus  we  have 
incidentally  observed  many  ways  in  which  acids,  bases,  and  salts  may 
be  prepared  that  do  not  involve  a  union  of  the  constituent  ions  and 
are  probably  not  ionic. 

Acids  and  Bases.  —  Oxygen  acids  can  almost  all  be  prepared 
from  the  anhydrides,  which  are  not  ionogens,  and  water.  Phosphoric 
acid,  sulphurous  acid  (p.  71),  hypochlorous  acid  (p.  268),  and  many 
other  acids  are  so  formed.  The  hydrogen  halides  are  all  producible  by 
union  of  the  constituent  elements.  Many  acids  are  formed  from  others 
when  the  latter  are  heated ;  for  example,  chloric  acid  from  hypochlo- 
rous acid  and  phosphoric  from  phosphorous  acid  (q.v.). 

Bases  are  formed  by  the  union  of  oxides  of  metals  with  water 
(p.  119). 

Salts.  —  The  dry  ways  of  forming  salts  are  very  numerous.  Thus, 
many  are  formed  by  direct  union  of  the  elements,  as  in  the  case  of 
chlorides  (p.  175),  sulphides  (p.  12),  and  other  simple  salts.  Many  are 
made  by  reduction  or  oxidation  from  other  salts,  as  potassium  chloride 
from  potassium  chlorate  (p.  65),  or  potassium  perchlorate  from  the  latter 
(p.  275).  Often  a  reducing  or  an  oxidizing  agent  is  used,  as  in  making 
potassium  nitrite  (q.v.)  from  the  nitrate,  and  lead  sulphate  from  lead 
sulphide  (q.v.).  Almost  all  oxygen  salts  can  be  obtained  by  the  union 
of  two  oxides,  as  calcium  carbonate  (q.v.)  from  calcium  oxide  and 
carbon  dioxide.  Ammonium  salts  are  formed  by  combination  of  am- 
monia, which  is  not  an  ionogen,  with  acids  (p.  183). 


THE   CHEMICAL   BEHAVIOR   OF   IONIC   SUBSTANCES          365 

In  manufacturing  salts,  methods  like  the  above,  as  well  as  those 
involving  ionic  actions,  are  very  commonly  used.  In  each  case  the 
cheapest  and  most  easily  accessible  materials  are  chosen,  and  the  least 
expensive  operation  is  selected. 

Neutralization  is  theoretically  the  simplest  ionic  way  of  getting  a  salt,  because 
the  water  can  be  removed  by  mere  evaporation.  Yet  most  of  the  salts  which  are 
on  the  market  are  made  by  the  use  of  other  actions.  In  fact*  the  pure  bases  and 
acids  are  usually  too  expensive  to  be  utilized  as  sources  of  salts. 

The  commonest  definition  of  a  salt,  as  a  substance  formed  by  the  neutraliza- 
tion of  an  acid  by  a  base,  is  open  to  many  objections.  It  is  logically  defective 
because  it  does  not  describe  what  a  salt  is,  but  one  method  of  making  a  salt, 
which  is  an  entirely  different  matter.  It  is  unfortunate  in  its  choice  amongst  pos- 
sible paralogisms,  because  neutralization  is  more  significant  as  a  method  of  forming 
water  than  as  a  means  of  preparing  a  salt.  And  finally,  as  we  have  just  seen,  the 
definition  has  not  even  the  excuse  of  practical  value,  for  most  salts  are  manu- 
factured by  entirely  different  reactions. 

Exercises.  —  1.  Using  the  data  in  regard  to  ionization  (p.  330), 
formulate  other  dissociations  according  to  the  models  on  p.  335. 

2.  Explain  fully  (cf.  p.  297)  the  effect  of  potassium  bromide  upon 
the  ionization  of  cupric  bromide  (p.  335). 

3.  Give  a  list  of  all  the  colorless  ionic  substances  you  can  think  of 
(p.  337). 

4.  Formulate  fully,  according  to  the  diagram  on  p.  340,  the  precip- 
itation of  barium  sulphate  (p.  342),  of  silicic  acid  from  sodium  silicate 
(p.  350),  of  zinc  hydroxide  from  zinc  sulphate  (p.  350),  of  silver  chlo- 
ride from  silver  sulphate,  and  the  liberation  of  hydrogen  chloride  by 
phosphoric  acid  (p.  179). 

5.  Give  a  list   of   the  specific  physical  and  chemical   properties 
(p.  341)  of  iodidion  and  of  hydrion. 

6.  Formulate   (p.  361)   the    displacement    of  iodine   by  bromine 
(p.  236),  and  of  bromine  by  chlorine  (p.  232). 

7.  Explain  the  acid  reaction  of  ferric  chloride  (FeCl3)  solution 
(p.  344). 

8.  Name  all  the  physical  components  in  aqueous  solutions  of  po- 
tassium hydroxide,  hydrogen  chloride,  and  sulphuric  acid  (cf.  p.  322). 

9.  Name  the  anions  and  cations  whose  formulas  are  used  on  p.  337. 

10.  Formulate  (p.  361)  the  actions  of  iron  and  of  aluminium  on 
dilute  hydrochloric  acid. 

11.  What  is  the  molar  concentration  (p.  250)  of  hydrion  in  N/10 
hydrogen   sulphide  (p.  330)   and  in  N/10  acetic  acid,  of  natrion  in 
N/2  sodium  chloride,  and  of  cuprion  in  N  cupric  nitrate  ? 


366  INORGANIC  CHEMISTRY 

12.  Combining  the  models  on  pp.  349,  354,  and  356,  formulate  the 
action  of  hydrochloric   acid   on  magnesium   hydroxide  and  on  zinc 
hydroxide. 

13.  Formulate  (p.  338)  and  discuss  the  action  of  sulphuric  acid 
upon  potassium  permanganate  (p.  306)  and  upon   potassium  dichro- 
mate  (p.  306). 

14.  Formulate  (p.  354)  the  neutralizations  mentioned  on  p.  357. 

15.  What  do  we  infer  (p.  359)  from  the  fact  that  the  solution  of 
sodium  hydrogen  sulphide  (NaHS)  is  neutral  ? 

16.  Can  you  invent  an  interaction  of  two  soluble  salts  in  which 
both  products  shall  be  insoluble? 

17.  To  which  classes  of  ionic  actions  do  those  of  iodine  on  hy- 
drogen sulphide  (p.  238),  and  of  magnesium  on  cold  water  (p.  97), 
belong  ?     Formulate  the  former  according  to  the  model  on  p.  342. 

18.  What  metals,  beside  platinum,  would  be  most  likely  to  form 
suitable  electrodes  for  an  electrolytic  cell  (p.  324)  ? 

19.  How  should  you  attempt  to  obtain  (p.  363)  a  pure  aqueous 
solution  of  the  acid  HAg(CN)2? 

20.  Formulate  (p.  363)  the  electrolysis  of  hydriodic  acid  and  that 
of  cupric  sulphate,  the  latter  between  copper  electrodes  (p.  325). 

21.  Give,  for  each  of  the  following,  two  definitions,  one  in  terms  of 
experimental  facts,  the  other  in  terms  of  the  ionic  hypothesis :  acid, 
base,  salt,  neutralization,  acid  salt,  mixed  salt. 


CHAPTER  XXI 
SULPHUR    AND    HYDROGEN    SULPHIDE 

Occurrence.  —  Free  sulphur  is  found  in  volcanic  regions,  where  it 
is  mixed  with  gypsum  and  other  minerals  and  occupies  the  pores  of 
pumice-stone.  Kocky  materials  accompanying  a  mineral  in  this  way 
are  called  the  matrix.  Here  and  there  non-volcanic  deposits,  formed  by 
the  action  of  bacteria,  have  been  met  with,  as  in  Louisiana  and  in 
Germany.  There  are  many  minerals,  compounds  containing  sulphur, 
which  are  chiefly  important  on  account  of  their  other  constituents. 
Sulphides  of  metals,  such  as  pyrite  (FeS2),  copper  pyrites  (CuFeS2), 
galena  (PbS),  zinc-blende  (ZnS),  and  sulphates,  like  gypsum  (CaS04, 
2H20),  barite  (BaS04),  and  celestite  (SrS04),  are  fairly  plentiful.  Sul- 
phur is  a  constituent  of  albumin  and  other  substances  found  in  the 
animal  body. 

Manufacture.  —  Most  sulphur  is  obtained  by  the  simple  process  of 
melting  it  away  from  the  accompanying  volcanic  rock  at  a  low  tempera- 
ture. The  liquid  sulphur  is  allowed  to  run  into  wooden  molds,  in 
which  it  solidifies  in  the  form  of  roll  sulphur.  For  many  purposes  this 
sulphur  is  sufficiently  pure.  To  produce  the  best  quality  it  is  sub- 
jected to  distillation  from  earthenware  retorts.  The  vapor  passes  into 
a  large  brick  chamber  and  condenses  upon  the  walls  and  floor  in  the 
form  of  a  fine  powder,  sold  as  flowers  of  sulphur.  When  the  chamber 
has  become  heated,  the  sulphur  condenses  in  the  form  of  a  liquid, 
which  is  drawn  off  and  cast  in  molds  as  before. 

The  greater  part  of  the  sulphur  of  commerce  comes  from  Sicily, 
where,  in  1898,  447,000  tons  were  manufactured  against  41,000  tons 
elsewhere.  A  certain  amount  is  obtained  in  Japan.  It  is  found  in 
California,  Nevada,  and  Louisiana.  Within  the  last  two  years  the 
product  from  the  last-named  source  has  superseded  Sicilian  sulphur  in 
the  American  market.  Sulphur  is  popularly  known  as  brimstone. 

Some  sulphur  is  also  made  from  poly  sulphides  (q.v.)  (which  give 
precipitated  sulphur),  from  the  waste  products  of  alkali  manufacture 
by  Chance's  process  (q.v.),  and  from  the  exhausted  material  used  in 
removing  sulphur  from  illuminating-gas  during  its  purification. 

367 


368  INORGANIC   CHEMISTRY 

Physical  Properties.  —  The  chief  physical  peculiarity  of  sulphur 
is  that,  instead  of  existing  in  three  physical  states  only,  like  water,  it 
possesses  two  familiar  and  perfectly  distinct  solid  forms  and  two  dif- 
ferent liquid  states  of  aggregation. 

1.  Native  sulphur  is  yellow,  has  a  sp.  gr.  2.06  and  melts  at  114.5°. 
It  is  almost  insoluble  in  water,  but  dissolves  freely  in  carbon  disul- 
phide  (40  parts  in  100  at  18°)  and  in  sulphur  monochloride  (q.v.). 
When   good  crystals   are  found,  they  belong  to  the  rhombic  system 
(Fig.  1,  p.  11).     The  solid  sulphur  obtained  by  evaporating  a  solu- 
tion shows,  as  a  rule,  more  perfect  crystalline  forms  than  does  native 
sulphur,  but  in  all  other  respects  is  identical  with  it.     Roll  sulphur  is 
the  same  substance  as  these  two,  although  the  crystals  in  their  growth 
have  interfered  with  one  another,  and  the  mass  is  crystalline,  simply, 
and   not  well  crystallized.     This   variety   is   called,    from  its  form, 
rhombic  sulphur. 

2.  When  a  large  mass  of  melted  sulphur  solidifies  slowly,  and  the 
crust  is  pierced  and  the  remaining  liquid  poured  out  before  the  whole 
has  become  solid,  the  interior  is  found  to  be  lined  with  long  needles. 
This  kind  of  sulphur  is  nearly  colorless  and  has  a  sp.  gr.  1.96,  melts 
at  119°,  and  is  in  all  physical  respects  a  different   individual  from 
rhombic  sulphur.     The  crystals  are  long,  transparent  prisms,  almost 
rectangular  in  section,  and  usually  showing  one  large  oblique  face  at 
the  end.     This  variety  of  the  simple  substance  is  named,  from  the 
system  to  which  its  crystals  belong,  monoclinic  sulphur. 

Monoclinic  sulphur  can  be  kept  only  above  96°  and  below  its  melting- 
point  (119°).  Every  recently  solidified  mass  of  sulphur  is  composed  of 
it.  But,  below  96°  the  mass  gradually  becomes  opaque,  the  change 
usually  spreading  from  one  or  two  points  and  finally  affecting  the 
whole  mass.  The  opacity  is  due  to  the  fact  that  the  material  has 
turned  into  an  aggregate  of  small  particles  of  rhombic  sulphur,  each 
of  which  occupies  less  space  than  the  monoclinic  sulphur  from  which 
it  was  formed.  Conversely,  rhombic  sulphur  can  be  kept  only  below 
96°.  When  heated  above  this  temperature,  but  not  as  high  as  its 
melting-point,  it  turns  slowly  into  monoclinic  sulphur.  Contact  with 
a  piece  of  monoclinic  sulphur,  or  mere  rubbing  with  a  hard  body,  will 
determine  the  point  at  which  the  transformation  shall  begin,  and  the 
expansion  which  accompanies  this  results  in  a  spreading  opacity  as 
before.  The  delay  before  the  change  starts  and  the  effect  of  rubbing 
and  inoculation  are  familiar  in  connection  with  almost  all  changes  of 
state  (cf.  p.  159). 


SULPHUR   AND  HYDROGEN   SULPHIDE  369 

Transitions,  marked  by  definite  points,  like  this  one  at  96°,  are 
attended  by  similar  phenomena,  whether  they  lie  between  two  solid 
states,  or  a  liquid  and  a  solid  state  (ice  and  water),  or  a  gaseous  and  a 
liquid  state  (steam  and  water).  Heat  is  given  out  when  we  pass  in  one 
direction  and  absorbed  by  passage  in  the  other.  The  rate  of  change 
of  vapor  pressure  with  change  in  temperature  (cf.  p.  163)  is  different 
on  each  side  of  the  transition  point.  A  body  which  has  two  solid  states, 
and,  therefore,  two  crystalline  forms,  is  said  to  be  dimorphous  (two- 
formed),  and  one  with  more  than  two  such  states  polymorphous  (see 
Ammonium  nitrate).  But  this  term  is  not  intended  to  imply  that  the 
relation  of  two  solid  states  to  each  other  is  essentially  different 
from  that  of  two  states  of  different  kinds,  such  as  solid  and  liquid, 
although  the  term  "  polymorphous  "  is  not  applied  to  the  latter. 

3.  When  melted  sulphur  is  heated,  it  undergoes  another  change  at 
or  above  160°.  The  pale-yellow,  mobile  liquid  (SA)  suddenly  becomes 
dark  in  color  and  so  viscous  (SM)  that  the  vessel  may  be  inverted 
without  loss  of  material.  Beyond  260°  the  viscidity  becomes  noticeably 
less,  and  at  445°  the  liquid  boils  and  passes  into  sulphur  vapor. 

The  sulphur  thermometer  thus  shows  many  more  fixed  points  than 
does  the  water  thermometer.  The  latter  consists  of  ice  up  to  0°,  of 
water  from  0°-100°,  and  of  steam  above  100°.  The  former  includes 
rhombic  sulphur  up  to  96°,  monoclinic  sulphur  96°-119°,  a  further, 
less  sharp  change  from  mobile  liquid  to  viscous  liquid,  and  then  vapor 
beyond  445°. 

Insoluble,  Amorphous  Sulphur. —  When  sulphur  which  has 
been  exposed  to  the  air,  and  particularly  sulphur  to  which  a  trace  of 
iodine  has  been  added,  is  boiled  and  then  allowed  slowly  to  cool,  the 
product  is  crystalline  and  soluble  in  carbon  disulphide,  as  before. 
But  when  such  impure  sulphur  is  boiled  and  then  suddenly  chilled  by 
pouring  into  cold  water,  it  is  at  first  semi-fluid.  After  several  days 
this  elastic  sulphur,  as  it  is  called,  becomes  hard.  It  is  then  found 
to  contain  rhombic  sulphur  mixed  with  a  large  proportion  of  another 
variety  of  free  sulphur.  This  is  almost  insoluble  in  any  solvent,  and 
so  may  be  secured  by  washing  the  mixture  with  carbon  disulphide. 

This  insoluble  sulphur,  being  without  crystalline  structure,  is  called 
also  amorphous  (Gk.  d  priv.,  nop<f>rj  form)  sulphur.  Now  amorphous 
bodies  (see  Glass)  are  always  supercooled  liquids,  that  is,  liquids  still 
existing  as  such  at  a  temperature  at  which  the  solid,  crystalline  form 
is  the  stable  one.  They  have  been  brought,  by  cooling,  so  rapidly 


370  INORGANIC  CHEMISTRY 

• 

through  their  freezing-point  (in  general,  transition  point  from  one 
physical  state  to  another)  that  crystallization  has  not  had  time  to 
begin  (cf.  p.  144)  and  a  general  rigidity  only  has  supervened.  Now 
amorphous  sulphur  is  viscous,  liquid  sulphur  (SM)  which,  by  sudden 
chilling,  has  been  carried  past  both  the  transition  to  mobile  liquid 
sulphur  (SA)  at  160°,  and  the  crystallization  as  well,  without  under- 
going either  of  these  changes.  It  is  supercooled  SM.  This  accounts 
for  the  fact  that  it  is  obtainable  only  by  rapid  cooling.  As  much  as 
34  per  cent  of  the  whole  may  thus  be  supercooled.  Once  the  mixture 
has  been  obtained  by  chilling,  the  insoluble  sulphur  reverts  very  slowly 
to  the  soluble  variety,  and  years  are  required  for  the  completion  of 
the  reversion  at  ordinary  temperature.  At  100°  the  reversion  is  com- 
pleted in  an  hour.  The  capacity  of  the  SM  to  be  supercooled  at  all 
seems  to  depend  on  the  presence  of  traces  of  foreign  bodies.  Of  these, 
iodine  is  the  most  efficient.  The  sulphur  dioxide  or  sulphuric  acid 
produced  by  prolonged  exposure  to  the  air  (see  below)  is  the  agent 
usually  responsible  for  the  supercooling.  Freshly  recrystallized  sul- 
phur gives  no  elastic  sulphur  and  no  insoluble  sulphur. 

Insoluble  sulphur  is  found  in  flowers  of  sulphur  and  in  sulphur 
formed  by  precipitation  from  thiosulphates  (q.v.)  in  presence  of  acids. 

Chemical  Properties.  —  When  the  density  of  sulphur  vapor  is 
determined  at  low  temperatures  and  under  reduced  pressures,  the 
molecular  weight  corresponds  closely  to  the  formula  S8.  As  the  tem- 
perature is  raised,  however,  the  vapor  expands  very  rapidly,  and  at 
800°  the  molecular  weight  is  64.2,  and  the  formula  therefore  S2. 
Intermediately  mixtures  of  S8,  S6,  S4,  and  S2  exist.  At  higher  tempera- 
tures no  further  dissociation  seems  to  take  place.  The  formula  of 
dissolved  sulphur,  as  measured  by  freezing-  and  boiling-point  methods 
(p.  162),  is  S8. 

We  do  not  ordinarily  think  of  sulphur  as  a  very  active  chemical 
substance,  but  this  is  largely  due  to  the  fact  that  its  solid  condition 
interferes  with  the  attainment  of  close  contact  with  the  body  upon 
which  it  acts.  When  finely  divided  metals,  with  the  exception  of  gold 
and  platinum  (cf.  p.  362),  are  rubbed  together  with  powdered  sul- 
phur, union  takes  place  and  sulphides  are  produced.  We  have  seen 
that  sulphur  when  heated  combines  with  great  vigor  with  iron  and 
copper,  as  it  does  indeed  with  most  of  the  metals.  Sulphur  unites 
also  with  many  of  the  non-metals.  Thus  with  oxygen  it  produces 
sulphur  dioxide,  and  even  sulphur  trioxide  (S0g).  It  unites  also  with 


SULPHUR  AND  HYDROGEN   SULPHIDE  371 

chlorine  directly.  When  sulphur  is  treated  with  oxidizing  agents  in 
presence  of  water,  no  trace  of  sulphur  dioxide  (or  sulphurous  acid) 
is  formed  ;  the  only  product  is  sulphuric  acid.  Even  free  oxygen  gas 
in  the  air  is  capable  slowly  of  oxidizing  moist  powdered  sulphur  and 
producing  sulphuric  acid,  2S  +  21^0  +  302  -»  2H2S04.* 

Uses  of  Sulphur.  —  Large  quantities  of  crude  sulphur  are  em- 
ployed for  making  sulphur  dioxide,  which  is  used  in  the  manufacture 
of  sulphuric  acid,  in  bleaching  feathers,  straw,  and  wool,  and  in 
making  alkali  sulphites  for  employment  in  the  bleaching  industry. 
The  manufacture  of  carbon  disulphide  consumes  a  considerable  amount 
also.  The  purified  sulphur  is  employed  in  the  manufacture  of  gun- 
powder, fireworks,  matches,  and,  by  combination  with  rubber,  of 
vulcanite.  Flowers  of  sulphur  is  used  in  vineyards  to  destroy  fungi, 
which  it  does  by  virtue  of  the  traces  of  sulphuric  acid  it  yields  by 
oxidation. 

HYDROGEN  SULPHIDE. 

This  compound  is  found  in  some  mineral  waters,  which  in  conse- 
quence are  known  as  sulphur  waters.  It  is  produced  in  the  decompo- 
sition of  animal  matter  containing  sulphur,  when  air  is  excluded,  and 
the  distinctive  odor  of  rotten  eggs  is  due  in  part  to  its  presence. 

Preparation.  —  1.  Hydrogen  and  sulphur  do  not  unite  percep- 
tibly in  the  cold.  At  310°  almost  complete  union  occurs,  but  about 
168  hours  are  required  for  the  change. 

2.  Sulphides  of  metals,  being  salts,  are  acted  upon  more  or  less 
easily  by  dilute  acids  (p.  349),  and  give  hydrogen  sulphide.  Ferrous 
sulphide,  the  least  expensive  of  those  easily  affected,  is  generally  used  : 

FeS  +  2HC1  ->  H2S  +  FeCL,. 

For  hydrochloric  acid  we  may  substitute  an  aqueous  solution  of  any 
active,  non-oxidizing  acid.  A  Kipp's  apparatus  (p.  97)  is  commonly 
employed.  Since  ferrous  sulphide  is  but  slightly  soluble  in  water,  the 
action  proceeds  by  a  rather  complex  series  of  equilibria  : 

FeS  (solid)  <=»  FeS  (diss'd)  <=>  Fe"  +  S" 


*  The  paragraph  on  the  chemical  relations  of  the  element  (see  end  of  this 
chapter)  should  be  read  at  this  point. 


372  INORGANIC  CHEMISTRY 

The  dissolved  hydrogen  sulphide  is  very  feebly  ionized,  and  main- 
tains a  smaller  concentration  of  snlphidiou  (S")  than  does  ferrous 
sulphide,  in  spite  of  the  comparative  insolubility  of  the  latter.  Hence, 
the  S"  formed  from  the  EeS  is  continuously  removed  by  union  with 
the  hydrion  furnished  by  the  acid,  S"-f-  2H*  £?  H2S,  and  all  the  other 
equilibria  are  constantly  displaced  forwards  on  this  account.  The 
action  is  therefore,  in  essence,  like  neutralization  (p.  354).  It  will  be 
observed  that  the  action  takes  place  rather  on  account  of  the  feeble 
ionization  of  the  weak  acid  than  by  reason  of  the  activity  of  the  other 
acid.  We  should  therefore  prefer  to  say  that  the  weak  acid  with- 
draws, and  not,  as  is  sometimes  done,  that  the  strong  acid  drives  it 
out. 

Since  the  action  is  an  ionic  one,  the  acids  must  be  employed  in 
dilute  form.  This  is  true  especially  of  oxygen  acids.  Thus,  concen- 
trated sulphuric  acid  has  little  action  upon  ferrous  sulphide  in  the 
cold,  and  when  the  substances  are  heated  the  oxygen  of  the  sulphuric 
acid  comes  into  play,  and  sulphur  dioxide  (q.v.)  and  free  sulphur  are 
formed. 

3.  Hydrogen  sulphide  is  the  invariable  product  of  the  extreme  re- 
duction of  any  sulphur  compound.  Thus,  it  is  formed  by  the  action 
of  hydrogen  iodide  upon  concentrated  sulphuric  acid  (p.  237).  Even 
sulphur  itself  is  reduced  by  dry,  gaseous  hydrogen  iodide : 

2HI  +  S  -»  H2S  +  I2. 

The  action  appears  to  be  just  the  reverse  of  that  which  takes  place  in 
aqueous  solution  (p.  238),  but  in  reality  is  quite  different.  Iodine  and 
gaseous  hydrogen  sulphide  will  not  produce  free  sulphur  and  gaseous 
hydrogen  iodide,  for  this  action  would  involve  a  considerable  increase 
in  energy  in  the  system.  But,  in  water,  they  do  give  hydrion  and 
iodidion,  for  these  bodies  contain  very  much  less  energy  than  does 
hydrogen  iodide : 

2H*  +  S"+I2-»2H-  +  S{+2I'    or     S"  +  I2  ->  S  J  +  21'. 

Physical  Properties.  —  Hydrogen  sulphide  is  a  colorless  gas 
with  a  characteristic  odor.  When  liquefied,  it  boils  at  —  60.4°  (765 
mm.),  and  in  solid  form  melts  at  —  82.9°.  At  12°  the  liquid  exerts  a 
pressure  of  15  atmospheres.  The  solubility  in  water  at  10°  is  360  vol- 
umes in  100,  and  becomes  less  as  the  temperature  is  raised.  The  gas 
can  be  driven  out  completely  by  boiling  the  solution  (cf.  p.  182).  The 


SULPHUR   AND   HYDROGEN   SULPHIDE  373 

gas  is  very  poisonous,  one  part  in  two  hundred  being  fatal  to  mammals, 
while  it  is  stated  that  one  part  in  fifteen  hundred  produces  death 
in  birds. 

Chemical  Properties  of  the  Gas.  —  When  heated,  the  gas  dis- 
sociates : 

H2S  ?±  H2  +  S, 

At  310°  the  decomposition,  although  very  slow,  affects  a  small  but 
perceptible  proportion  of  the  gas  before  coming  to  rest.  The  dissoci- 
ation, like  most  thermal  dissociations,  is  accompanied  "by  an  absorption 
of  heat  and  is  therefore  greater  at  higher  temperatures  (cf.  p.  260). 

Hydrogen  sulphide  forms  a  solid  hydrate  with  water,  H8S,  7H2O,  the  behavior 
of  which  resembles  that  of  chlorine  hydrate  (p.  174).  The  vapor  tension  of  water 
and  the  gas  amount,  together,  to  1  atmosphere  at  5°.  The  compound,  therefore, 
decomposes  rapidly  at  5°,  and  more  slowly  at  lower  temperatures. 

The  gas  burns  in  air,  forming  steam  and  sulphur  dioxide.  The 
temperature  of  the  mantle  of  flame  surrounding  the  gas,  as  it  issues 
from  a  jet,  being  far  above  310°,  the  gas  in  the  interior  is  dissociated 
before  it  meets  with  any  oxygen.  Hence  a  cold  dish  held  across  the 
flame  receives  a  deposit  of  free  sulphur,  and  a  part  of  the  hydrogen 
also  escapes  unburnt.  It  may  be  remarked  that  dissociation  of  this 
kind  probably  precedes  the  combustion  of  most  gaseous  compounds 
(see  Flame). 

The  metals,  down  to  and  including  silver  in  the  electromotive 
series,  when  exposed  to  the  gas,  quickly  receive  a  coating  of  sulphide. 
That  the  gas  should  thus  behave  like  free  sulphur  shows  its  instability. 

This  instability  is  shown  also  in  the  fact  that  its  hydrogen  reduces 
substances,  such  as  sulphur  dioxide,  which  are  not  affected  by  free 
hydrogen : 

2H2S  +  S02  -»  2H20  +  3S. 

This  action  takes  place  much  more  rapidly  when  the  gases  are  moist 
than  when  they  are  dry,  and  is  retarded  by  dilution  with  indifferent 
gases  (cf.  p.  252).  Native  sulphur  is  probably  produced  by  this  action, 
as  both  of  these  gases  are  found  issuing  from  the  ground  in  volcanic 
neighborhoods.  Sulphur  is  deposited  also  when  hydrogen  sulphide 
undergoes  a  partial  combustion  with  a  restricted  supply  of  oxygen  : 

2H2S  +  02  -*  2H20  +  28. 
Its  formation  in  nature  is  sometimes  to  be  accounted  for  in  this  way. 


374  INORGANIC   CHEMISTRY 

When  hydrogen  sulphide  gas  is  led  through  concentrated,  or  even, 
simply,  normal  sulphuric  acid,  the  acid  is  reduced,  sulphur  dioxide 
escapes,  and  sulphur  is  deposited  : 

H2S  +  H2S04  ->  S  +  2H20  +  SO2. 

The  sulphuric  acid  may  be  written  H20,  S08.  In  furnishing  S02,  there- 
fore, each  molecule  can  give  one  unit  of  oxygen  and  therefore  oxidize 
one  molecule  of  H2S  (see  p.  389).  On  account  of  this  action  the  gas 
cannot  be  dried  by  means  of  concentrated  sulphuric  acid.  Calcium 
chloride  is  likewise  inapplicable,  since  a  partial  interchange  takes 
place,  resulting  in  the  production  of  calcium  sulphide  and  hydrogen 
chloride  gas.  Only  a  dehydrating  agent,  such  as  phosphoric  anhy- 
dride, with  which  it  cannot  interact,  is  suitable  for  drying  the  gas. 

Chemical  Properties  of  the  Aqueous  Solution  of  Hydrogen 
Sulphide.  —  While  the  gas  itself  is  not  an  acid,  its  solution  in  water 
gives  a  feeble  acid  reaction  with  litmus.  In  an  aqueous  solution 
(N/10),  only  .0007  (0.07  per  cent)  of  the  substance  is  ionized  accord- 
ing to  the  equation  (of.  p.  346) : 

H2S  ->  H-  +  HS'. 

Some  S"  ions  are  present,  but  hydrosulphidion  (HS')  is  less  dissoci- 
ated than  is  water  itself,  and  the  amount  of  sulphidion  is  therefore 
very  small. 

By  the  action  of  oxygen  from  the  air  upon  an  aqueous  solution  of 
hydrogen  sulphide,  the  sulphur  is  slowly  displaced  and  appears  in  the 
form  of  a  fine  white  powder : 

O2  +  2H2S  ->  2S  +  2H20. 

This  is  an  action  similar  to  the  displacement  of  ionic  iodine  by  free 
chlorine  (p.  236).  On  the  other  hand,  the  hydrogen  may  be  displaced 
by  metals,  particularly  the  more  active  ones,  but  the  small  degree  of 
ionization  makes  the  action  very  slow. 

The  solution  of  the  gas  is  a  reducing  agent,  as  its  action  upon 
iodine  shows  (p.  238).  So,  also,  in  presence  of  an  acid,  it  removes 
oxygen  from  dichromic  acid  (produced  by  the  action  of  an  acid  upon 
potassium  dichromate) : 

K2Cr207    +    2HC1  <r»  H2Cr207  +  2KC1  (1) 

H2Cr207    +    6HC1  ->  4H20  +  2CrCl8(+  30)  (2) 

(30)         +    3H2S  -»  3H20  +  3S  (3) 

Adding:  K2Cr2O7+8HCl+3H2S  -»  2KC1  +  2CrCl8+7H20  +  3S 


SULPHUR   AND   HYDROGEN   SULPHIDE  375 

The  first  partial  equation  (cf.  p.  228)  represents  the  regular  interaction 
of  two  ionogens,  but  the  second  interaction  does  not  take  place  unless  an 
oxidizable  body  (here  the  hydrogen  sulphide)  is  present  to  take  posses- 
sion of  the  oxygen  which  it  is  capable  of  delivering  (cf.  p.  307). 

As  an  acid,  the  solution  of  hydrogen  sulphide,  sometimes  known  as 
sulphydric  acid,  may  be  neutralized  by  bases. 

Sulphides.  —  As  a  di-basic  acid  (p.  358),  hydrogen  sulphide  gives 
both  acid  and  neutral  (or  normal)  sulphides,  such  as  NaHS  and  Na,jS. 

The  former  are  obtained  by  passing  the  gas  in  excess  into  solutions 
of  soluble  bases  : 

H2S  +  NaOH  ->  H20  +  NaHS, 

and  are  neutral  in  reaction.  Their  negative  ion,  HS',  gives  practically 
no  hydrion. 

By  adding  to  the  above  solution  an  amount  of  sodium  hydroxide 
equal  to  that  used  before,  and  driving  off  the  water  by  evaporation, 
the  second  unit  of  hydrogen  is  displaced  : 

NaOH  +  NaHS  ±=5  Na2S  +  H20  f  . 

This  action  is  wholly  reversed  when  dry  sodium  sulphide  is  dissolved 
in  water,  the  salt  being  completely  hydrolyzed  to  the  acid  salt  : 


S"  )<_  Hg. 
H20  ±+  OH'  -f-  H*   \ 

The  HS'  gives  a  lower  concentration  of  hydrion  than  the  water,  and 
hence  uses  up  in  its  formation  the  ions  of  hydrogen  produced  by  the 
latter  until  an  amount  of  hydroxyl  equivalent  to  half  the  sodium  is 
formed. 

Many  sulphides  are  insoluble  in  water,  and  these  may  be  divided 
roughly  into  three  classes  : 

1.  The  sulphides  of  silver,  copper,  mercury,  and  some  other 
metals  are  exceedingly  insoluble,  and,  therefore,  do  not  interact  with 
dilute  acids  as  does  ferrous  sulphide  (p.  371).  These  may  therefore 
be  made  by  leading  hydrogen  sulphide  into  a  solution  of  any  of  their 
salts  : 

CuS|  +  H2S04. 


The  acid  produced  has  scarcely  any  effect  upon  the  sulphide,  and  al- 
most no  reverse  action  is  observed. 


376  INORGANIC   CHEMISTRY 

2.  The  sulphides  of  iron,  zinc,  and  certain  other  metals  are  insol- 
uble in  water,  but  not  so  much  so  as  the  last  class.     Hence  they  are 
decomposed  by  dilute  acids,  and  the  reverse  of  the  above  action  takes 
place  almost  completely.     These   sulphides  must  therefore  be  made, 
either  by  combination  of  the  elements,  or  by  adding  a  soluble  sulphide 
to  a  solution  of  a  salt  : 

FeS04  +  (NH4)2S  +=»  FeS  J  +  (NH4)2S04. 

No  acid  is  produced  in  this  sort  of  interaction,  and  the  considerable 
insolubility  of  the  sulphide  of  iron  or  zinc  in  water  renders  the 
change  nearly  complete. 

3.  The  sulphides  of  barium,  calcium,  and  some  other  metals  (q.v.), 
although  insoluble  in  water,  are  hydrolyzed  by  it,  and  give  soluble 
products  (see  p.  604)  : 

2CaS  +  2H20  <=;  Ca(OH)2  +  Ca(SH)2. 

They  may  be  prepared  by  direct  union  of  the  elements,  and  from  the 
sulphates  by  reduction  with  carbon.  (For  the  application  of  the  above 
differences  in  solubility  amongst  sulphides  in  chemical  analysis,  see 
end  of  Chap,  xxxvii.) 

The  soluble  acid  sulphides  are  oxidized  in  aqueous  solution  by  at- 
mospheric oxygen  : 

2NaSH  +  O2  -*  2NaOH  +  2S. 

The  sulphur  is  not  precipitated,  but  combines  with  the  excess  of  the 
sulphide,  forming  poly  sulphides  (see  below). 

Poly  sulphides.  —  When  sulphur  is  shaken  with  a  solution  of  an 
alkaline  sulphide  or  acid  sulphide,  it  dissolves,  and  evaporation  of  the 
solution  leaves  substances  varying  in  composition  from  Na^  to  Na^Sj. 
Whether  those  containing  less  than  five  units  of  sulphur  are  mixtures 
of  the  pentasulphide  with  the  ordinary  sulphide  (Na^S)  in  different 
proportions,  or  are  single  substances,  has  not  been  determined. 

When  an  acid  is  poured  into  sodium  pentasulphide  solution,  minute 
spherules  of  rhombic  sulphur  are  precipitated  : 

Na2S5  +  2HC1  ->  2NaCl  -f  H^Sf  +  48  1. 


When  the  order  is  reversed,  sodium  pentasulphide  being  thrown 
into  concentrated  hydrochloric  acid,  no  hydrogen  sulphide  is  evolved. 
Hydrogen  pentasulphide,  H2S5,  a  yellow  oil,  falls  to  the  bottom  of  the 
vessel. 


SULPHUR   AND   HYDROGEN   SULPHIDE  377 

When  purified,  the  compound  is  nearly  colorless.  Sulphur  dis- 
solves in  it  freely,  giving  a  yellow  solution.  It  can  be  kept  without 
change  when  dry,  but  contact  with  water  causes  it  to  decompose  into 
hydrogen  sulphide  and  sulphur. 

The  Chemical  Relations  of  the  Element.  —  In  combination  with 
metals  and  hydrogen,  sulphur  is  bivalent,  forming  compounds  like 
H2S,  FeS,  CuS,  and  HgS.  In  combination  with  non-metals,  however, 
the  valence  is  frequently  greater,  the  maximum  being  seen  in  sulphur 
trioxide,  where  we  must  assume  that  the  sulphur  is  sexivalent.  Its 
oxides  are  acid-forming,  and  it  is,  therefore,  a  non-metal. 

Sulphur  is  regarded  as  resembling  oxygen  more  closely  than  any 
of  the  other  elements  we  have  studied  so  far.  Both  unite  directly  with 
most  metals  and  non-metals.  In  this  they  are  like  chlorine.  But  hy- 
drogen chloride  is  highly  ionized  by  water,  while  the  hydrogen  com- 
pounds of  oxygen  and  sulphur  are  feebly  ionized.  The  formulae  of  the 
compounds  of  oxygen  and  sulphur  with  metals  are  similar,  CuO  and 
CuS,  NaOH  and  NaSH,  and  so  forth,  but  this  is  in  part  due  merely  to 
the  fact  that  both  elements  are  bivalent.  The  chemical  resemblance  of 
sulphur  to  selenium  and  tellurium  (q.v.)  is  much  more  striking  than 
its  resemblance  to  oxygen. 

Exercises.  —  1.  The  freezing-point  of  pure  sulphur  is  found  to 
vary  from  119°  down  to  114°,  depending  upon  the  temperature  to 
which  the  liquid  has  been  heated  and  the  speed  with  which  it  has  been 
cooled.  To  what  should  you  suspect  this  variability  to  be  attributable  ? 

2.  How  could  the  decomposition  of  hydrogen  sulphide  at  310°  be 
rendered  (a)  more  complete,  (b)  less  complete  ?    Would  the  percentage 
decomposed  be  affected  (a)  by  reducing  the  pressure,  (#)  by  mixing 
the  gas  with  an  indifferent  gas  ? 

3.  To  what  classes  of  ionic  actions  (p.  360)  do  the  interactions  of 
hydrogen    sulphide   solution    and    (a)    oxygen  (p.  374),  (b)  acidified 
potassium   dichromate   (p.  374),   (c)   sodium  hydroxide  (p.  375),  (d) 
iodine  (p.  372),  belong  ? 

4.  Why  is  normal  sodium  sulphide  only  half  hydrolyzed  by  water  ? 


CHAPTER    XXII 
THE    OXIDES    AND    OXYGEN    ACIDS    OF    SULPHUR 

FOUR  oxides  of  sulphur,  represented  by  the  formulae  S203,  S02, 
S03,  and  S207,  are  known.  Of  these,  however,  the  first  and  the  last 
are  much  less  familiar  substances  than  the  other  two.  The  dioxide 
and  trioxide  of  sulphur  are  not  only  important  in  themselves,  but 
their  relation  to  the  acids  H2S03  and  H2S04,  which  may  be  obtained 
from  them  by  the  addition  of  water,  makes  them  doubly  so  to  the 
chemist. 

Preparation  of  Sulphur  Dioxide.  —  When  sulphur  burns  in  air 
or  oxygen,  sulphur  dioxide  is  produced  (p.  67).  While  some  of  the 
sulphur  dioxide  used  in  commerce  is  prepared  in  this  manner,  the 
larger  part  is  probably  obtained  by  the  roasting  of  sulphur  ores. 
Pyrite  (FeS2),  for  example,  which  is  a  familiar  yellow,  metallic-look- 
ing mineral,  burns  when  it  has  been  raised  to  the  kindling  tempera- 
ture, on  account  of  the  large  amount  of  sulphur  which  it  contains  : 

4FeS2  +  1102  ->  2Fe208  +  8SO2. 

It  should  be  noted  in  passing,  that  heating  and  roasting  are  dis- 
tinct processes  in  chemistry.  The  latter  term  always  assumes  the 
access  of  the  air  and  employment  of  its  oxygen  ;  the  former,  in  the 
absence  of  modifying  words,  assumes  the  exclusion  or  the  chemical  in- 
difference of  the  air. 

Sulphur  dioxide  is  also  set  free  by  the  action  of  acids  upon 
sulphites.  Sulphuric  acid  and  a  strong  solution  of  sodium  sulphite 
may  be  used  : 

" 


<'    >'       so> 

The  sulphurous  acid  being  only  moderately  ionized,  its  molecules  are 
formed  in  considerable  amount.  Being  also  very  unstable,  it  decom- 
poses spontaneously  into  water  and  sulphur  dioxide,  and  the  latter 
escapes  when  sufficient  water  for  its  solution  is  not  present. 

378 


THE   OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR  379 

In  the  laboratory,  sulphur  dioxide  is  frequently  made  by  the 
reduction  of  concentrated  sulphuric  acid.  Copper  is  the  metal  most 
commonly  employed,  because  by  its  means  very  pure  sulphur  dioxide 
can  be  obtained.  More  active  metals,  such  as  iron  and  zinc,  although 
cheaper,  cannot  be  used,  since  they  reduce  the  sulphuric  acid  to 
hydrogen  sulphide.  The  undiluted  hydrogen  sulphate  consists  en- 
tirely of  molecules,  and,  at  the  high  temperatures  at  which  alone  the 
action  is  vigorous,  is  an  oxidizing  agent  (cf.  p.  97).  A  part  of  it  loses 
oxygen  to  form  water  with  the  hydrogen  of  another  molecule: 


H2S04-»H20  +  S02(  +  0)  (1) 

(0)  +  H2S04  +  Cu  -»  H20  +  CuS04  _  (2) 

2H2S04  +  Cu  -»  2H20  +  S02  +  CuS04 

Some  easily  oxidized  non-metals,  such  as  carbon  and  sulphur,  act  in 
the  same  way  : 

C  +  2H2S04  ->  2H20  +  2S02  +  C02, 

S  -fc  2H2S04  -»  2H20  +  3S02. 

Physical  Properties.  —  Sulphur  dioxide  is  a  gas  possessing  a 
penetrating  and  characteristic  odor.  This  is  frequently  spoken  of  as 
the  "  odor  of  sulphur,"  but  it  should  be  remembered  that  sulphur 
itself  has  scarcely  any  smell  at  all.  The  weight  of  the  G-.M.V.  of 
the  gas  (65.55  g.)  shows  it  to  be  more  than  twice  as  heavy  as  air. 
The  critical  temperature  is  156°.  By  means  of  a  freezing  mixture  of 
ice  and  salt,  the  gas  is  easily  condensed  to  a  transparent,  mobile  fluid, 
which  boils  at  —  8°.  At  20°  the  vapor  tension  of  the  liquid  is  3.25 
atmospheres.  The  liquid  may  be  frozen  to  a  white  solid,  melting  at 
—  76°.  It  ionizes  substances  dissolved  in  it  as  well  as  does  water.  The 
solubility  of  the  gas  in  water,  5000  volumes  in  100,  is  very  great. 
Unlike  solutions  of  the  hydrogen  halides  (p.  182),  however,  the  liquid 
is  completely  freed  from  the  gas  by  boiling. 

Chemical  Properties.  —  Sulphur  dioxide  is  stable,  being  decom- 
posed only  by  the  use  of  a  very  high  temperature. 

It  unites  with  water  to  form  sulphurous  acid,  H2SOS.  Although 
the  gas  itself  sometimes  receives  this  name,  it  is  not  acid  :  it  is  simply 
the  anhydride  (p.  71)  of  the  acid. 

Since  the  maximum  valence  of  sulphur  is  6,  sulphur  dioxide,  in 
which  but  four  of  the  valences  of  sulphur  are  used,  is  unaaturated. 
It  is  therefore  still  able  to  combine  directly  with  suitable  elements, 


380  INORGANIC   CHEMISTRY 

such  as  chlorine  and  oxygen.     When  it  is  mixed  with  chlorine  in  sun- 
light, a  liquid,  sulphuryl  chloride  S02CLj  is  produced. 

Liquefied  sulphur  dioxide  is  now  sold  in  tin  cans,  and  is  employed 
for  bleaching  straw,  wool,  and  silk.  As  a  disinfectant  it  has  been 
displaced  to  a  large  extent  by  formaldehyde. 

Preparation  of  Sulphur  Trioxide. —  Although  the  formation  of 
sulphur  trioxide  S03  is  accompanied  by  the  liberation  of  much  heat, 
sulphur  dioxide  and  oxygen,  even  when  heated  together,  unite  very 
slowly.  Ozone,  however,  combines  with  the  former  readily. 

The  interaction  of  sulphur  dioxide  and  oxygen  is  hastened  by 
many  substances,  such  as  glass,  porcelain,  ferric  oxide,  and,  more 
especially,  finely  divided  platinum,  which  remain  themselves  un- 
changed and  simply  act  as  catalytic  agents.  The  "contact  process," 
as  this  is  called,  has  been  rendered  available  for  the  commercial 
manufacture  of  sulphur  trioxide  by  Knietsch  (1901).  The  chief 
features  of  the  process  are  :  (1)  The  complete  removal  of  arsenious 
oxide  and  other  impurities  derived  from  the  burning  of  crude  pyrite, 
the  minutest  traces  of  which  poison  the  catalytic  agent  and  soon 
render  it  absolutely  inoperative.  (2)  The  preliminary  passage  of  the 
cold  mixture  of  gases  over  the  outside  of  the  pipes  containing  the 
contact  agent.  This  removes  part  of  the  heat  generated  by  the  ac- 
tion, S02  +  0  <=±  S08  +  22,600  cal.,  going  on  inside,  and  keeps  the 
temperature  of  the  interior  at  400°.  Below  400°  the  union  is  too 
slow  ;  above  400°,  being  reversible,  it  is  incomplete.  At  400°,  98-99 
per  cent  of  the  materials  unite ;  at  700°,  only  60  per  cent,  at  900°  none. 
Twice  the  quantity  of  oxygen  theoretically  needed  is  employed.  The 
vaporous  product  is  condensed  by  being  led  into  97-99  per  cent 
sulphuric  acid,  and  the  concentration  of  the  liquid  is  constantly 
maintained  at  this  point  by  the  regulated  influx  of  water.  The 
trioxide  is  thus  chiefly  used  for  immediate  conversion  into  sulphuric 
acid. 

Formerly  sulphur  trioxide  was  obtained  by  the  distillation  of 
impure  ferric  sulphate,  Fe2(S04)3  — >  Fe203  +  3S08.  It  may  also  be 
prepared  by  repeated  distillation  of  concentrated  sulphuric  acid  with 
a  powerful  drying  agent,  like  phosphoric  anhydride. 

Physical  Properties.  —  Sulphur  trioxide  S03  is,  at  ordinary  tem- 
peratures, fluid.  The  crystals,  obtained  by  cooling,  melt  at  14.8°. 
The  liquid  boils  at  46°,  and  is,  therefore,  exceedingly  volatile  at  ordi- 


THE   OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR  381 

nary  temperature.  It  fumes  strongly  when  exposed  to  the  air,  in  con- 
sequence of  the  union  of  the  vapor  with  moisture  and  the  production 
of  minute  drops  of  sulphuric  acid. 

A  white  crystalline  variety  of  the  substance,  which  in  appearance 
closely  resembles  asbestos,  is  obtained  when  a  trace  of  water  has 
gained  access  to  the  oxide.  It  is  a  polymer  (p.  242)  of  the  above 
liquid,  having  twice  its  molecular  weight,  and  therefore  the  formula 
(S08)2.  When  heated  at  50°  this  polymer  begins  to  dissociate  and 
passes  into  vapor  of  S08  without  melting.  This  white  solid  is  the 
more  stable  and  more  familiar  form  of  the  trioxide. 

Chemical  Properties. —  The  vapor  of  sulphur  trioxide  disso- 
ciates above  400°  into  sulphur  dioxide  and  oxygen. 

Sulphur  trioxide  is  not  itself  an  acid,  but  it  is  the  anhydride  of 
sulphuric  acid.  When  placed  in  water  it  unites  vigorously,  causing 
a  hissing  noise  due  to  the  steam  produced  by  the  heat  of  the  union. 
In  consequence  of  its  great  tendency  to  combine  w'ith  water,  the  liquid 
variety,  which  is  the  more  active,  removes  the  elements  of  this  sub- 
stance from  materials  which  contain  them  in  the  proper  proportions. 
Thus  paper,  which  is  largely  cellulose,  (CeH^Og).,.,  and  sugar,  C12H220U, 
are  charred  by  it,  and  carbon  is  set  free. 

Just  as  sulphur  trioxide  unites  with  water  to  give  hydrogen  sul- 
phate, so  it  combines  vigorously  with  many  oxides  of  metals,  produ- 
cing the  corresponding  sulphates : 

H20  +  SOS  fc?  H2S04 ,  CaO + SO8  -»  CaS04 . 

The  union  of  an  oxide  of  a  non-metal  with  the  oxide  of  a  metal,  in 
this  fashion,  is  a  general  method  of  obtaining  salts  (cf.  p.  364). 

Oxygen  Acids  of  Sulphur. —  Sulphurous  and  sulphuric  acids 
have  been  mentioned  frequently  already.  Next  to  them  in  importance 
come  thiosulphuric  acid  and  persulphuric  acid.  The  compositions  of 
the  acids  show  their  relationships  : 

Hyposulphurous  acid,  H2S204        Thiosulphuric  acid,  H2S2O8 
Sulphurous  acid,  H2S08          Persulphuric  acid,    H^Gg. 

Sulphuric  acid,  H^O* 

Thiosulphuric  acid  (Gk.  Oetov,  sulphur)  is  so  named  because  it  con- 
tains one  unit  of  sulphur  in  place  of  one  of  the  units  of  oxygen  of 
sulphuric  acid.  Besides  the  above  we  have  also :  Dithionic  acid 


382  INORGANIC   CHEMISTRY 

trithionic  acid  H2S8O6,  tetrathionic  acid  H2S406,  and  pentathionic  acid 
H&0r 

On  account  of  its  commercial  importance  and  the  interest  attach- 
ing to  its  method  of  manufacture  and  to  its  properties,  we  may  first 
discuss  sulphuric  acid.  We  shall  then  be  able  to  dispose  of  the  re- 
maining acids  in  a  much  briefer  fashion. 


SULPHURIC  ACID. 

Although  salts  of  sulphuric  acid,  such  as  calcium  sulphate,  are  ex- 
ceedingly plentiful  in  nature,  the  preparation  of  the  acid  by  chemi- 
cal action  upon  the  salts  is  not  practicable.  The  sulphates,  indeed, 
interact  with  all  acids,  but  the  actions  are  reversible.  The  comple- 
tion of  the  action  by  the  plan  used  in  making  hydrogen  chloride 
(p.  180),  involving  the  removal  of  the  sulphuric  acid  by  distillation, 
would  be  difficult  on  account  of  the  in  volatility  of  this  acid.  It  boils 
at  330°;  and  active  acids,  less  volatile  still,  which  might  be  used  to 
liberate  it,  do  not  exist.  We  are  therefore  compelled  to  build  up  sul- 
phuric acid  from  its  elements. 

The  union  of  sulphur  dioxide  and  oxygen  by  the  contact  process, 
and  combination  of  the  trioxide  with  water  (p.  380),  is  the  best  method 
for  making  a  highly  concentrated  acid.  For  obtaining  ordinary  "  oil 
of  vitriol,"  however,  an  entirely  different  plan  is  still  used  extensively. 

History  of  Sulphuric  Acid  Manufacture,  —  Impure  forms  of 
sulphuric  acid  have  been  known  for  many  centuries.  In  the  fifteenth 
century  it  was  made  by  distilling  ferrous  sulphate  with  sand.  The 
product,  however,  contained  much  water  and  sulphur  dioxide.  The 
first  successful  preparation  of  the  substance  commercially  was  made 
by  Ward  at  Eichmond-on-the-Thames  (1758).  The  process  consisted 
in  burning  a  mixture  of  sulphur  and  saltpeter  (KNO8)  in  a  ladle 
suspended  in  a  large  glass  globe  partially  filled  with  water.  The 
gases  which  were  evolved  contained  large  quantities  of  sulphur 
dioxide  and  oxides  of  nitrogen,  which  by  interaction  with  atmospheric 
oxygen  and  water  (see  below)  produced  the  sulphuric  acid.  The  solu- 
tion which  was  obtained,  although  it  could  be  prepared  of  any  desired 
concentration  by  the  burning  of  a  sufficient  number  of  charges,  was 
far  from  pure  and  was  expensive,  bringing  thirteen  shillings  ($3.25) 
per  pound.  Subsequently  a  chamber  lined  with  lead  was  substituted 


THE   OXIDES  AND  OXYGEN   ACIDS  OF  SULPHUR  383 

for  the  glass  vessel.  This  reduced  the  price  to  about  two  shillings 
and  sixpence  ($0.60)  per  pound.  The  same  principles  are  used  in 
the  modern  "  chamber  process." 

Chemistry  of  the  Chamber  Process.  —  The  gases,  the  interao 
tions  of  which  result  in  the  formation  of  sulphuric  acid,  are  :  water 
vapor,  sulphur  dioxide,  nitrous  anhydride  N2O8  (q.v.),  and  oxygen. 
These  are  obtained,  the  first  by  injection  of  steam,  the  second  usually 
by  the  burning  of  pyrite,  the  third  from  nitric  acid,  and  the  fourth  by 
the  introduction  of  air.  The  gases  are  thoroughly  mixed  in  large 
leaden  chambers,  and  the  sulphuric  acid  condenses  and  collects  upon  the 
floors.  In  spite  of  elaborate  investigations,  instigated  by  the  exten- 
sive scale  upon  which  the  manufacture  is  carried  on  and  the  immense 
financial  interests  involved,  some  uncertainty  still  exists  in  regard  to 
the  precise  nature  of  the  chemical  changes  which  take  place.  Accord- 
ing to  Lunge  the  greater  part  of  the  product  is  formed  by  two  succes- 
sive actions,  the  first  of  which  yields  a  complex  compound  that  is 
decomposed  by  excess  of  water  in  the  second  : 

H30  +  2S02  +  N,0.  +  02  -»  2S02<  ol"0  (1) 


The  group  —NO  is  found  in  many  compounds.  Here,  if  it  were  dis- 
placed by  hydrogen,  sulphuric  acid  would  result.  Hence  this  com- 
pound is  called  nitrosylsulphuric  acid: 


r\  _  TT  OTT 


2SO'0-NO          »    *          *     OH 

The  equations  (1)  and  (2)  are  not  partial  equations  for  one  inter- 
action, but  represent  distinct  actions  which  can  be  carried  out  sepa- 
rately. In  a  properly  operating  plant,  indeed,  the  nitrosylsulphuric 
acid  is  not  observed.  But  when  the  supply  of  water  is  deficient,  white 
"  chamber  crystals,"  consisting  of  this  substance,  collect  on  the  walls. 
The  explanation  of  the  success  of  this  seemingly  roundabout 
method  of  getting  sulphuric  acid  is  as  follows  :  The  direct  union  of 
sulphur  dioxide  and  water  to  form  sulphurous  acid  is  rapid,  but  the 
action  of  free  oxygen  upon  the  latter,  2H2S03  +  02  —  >  2HJS04,  is  ex- 
ceedingly slow.  Reaching  sulphuric  acid  by  the  use  of  these  two 
changes,  although  they  constitute  a  direct  route  to  the  result,  is  not 
feasible  in  practice.  On  the  other  hand,  both  of  the  above  actions, 
(1)  and  (2),  happen  to  be  much  more  speedy,  and  so,  by  their  use, 


384  INORGANIC   CHEMISTRY 

more  rapid  production  of  the  desired  substance  is  secured  at  the  ex- 
pense of  a  slight  complexity.  It  may  be  added  that  the  heat  finally 
given  out  in  the  formation  of  one  formula-weight  of  sulphuric  acid  is 
exactly  the  same  in  amount  whether  nitrous  anhydride  intervenes  or 
not  (cf.  p.  78). 

•  The  progress  of  the  first  action  is  marked  by  the  disappearance  of 
the  brown  nitrous  anhydride,  and,  on  the  introduction  of  water,  the 
completion  of  the  second  results  in  the  reproduction  of  the  same  sub- 
stance. It  would  thus  seem  as  if  the  nitrous  anhydride  should  take 
part  an  indefinite  number  of  times  in  these  changes  and  so  facilitate 
the  conversion  of  an  unlimited  amount  of  sulphur  dioxide,  oxygen, 
and  water  into  sulphuric  acid,  without  impairment  of  its  quantity. 
In  practice,  however,  certain  subsidiary  actions  take  place,  such  as, 
for  example,  the  reduction  of  some  nitrous  anhydride  to  nitrous  oxide 
(N20),  which  permanently  remove  a  part  of  the  material  from  parti- 
cipation in  the  cycle. 

The  supply  of  nitrous  anhydride  is  maintained  by  the  introduction 
of  nitric  acid  vapor  into  the  chamber.  This  acid  is  secured  by  the 
action  of  concentrated  sulphuric  acid  upon  commercial  sodium  nitrate 
NaN08: 

NaN08+H2S04  ?±  HN03|  +NaHSO4. 


On  account  of  the  volatility  of  the  nitric  acid,  a  moderate  heat  is 
sufficient  to  remove  it  from  admixture  with  the  other  substances,  and 
its  vapor  is  swept  along  with  the  other  gases  into  the  apparatus.  The 
initial  action  which  the  nitric  ^acid  undergoes  may  be  represented  by 
the  following  equation  : 


2S02  +  2HNO8  ->  2H2SO4  +  N208. 
We  may  write  this  in  the  form  : 

H20  -f  2S02  +  H20,N205  ->  211,80,  +  N203. 


The  two  molecules  of  water,  one  actually,  the  other  potentially,  pres- 
ent, with  the  two  molecules  of  sulphur  dioxide,  can  furnish  two 
molecules  of  sulphurous  acid  (H2S03).  The  N206  in  passing  to  the 
condition  N203  gives  up  the  two  units  of  oxygen  required  to  convert 
this  sulphurous  acid  into  sulphuric  acid. 

It  should  be  remarked  that  we  shall  later  find  nitrous  anhydride 
to  be  an  unstable  substance,  which,  in  the  gaseous  form,  is  composed 
largely  of  the  products  formed  by  its  dissociation,  N208  +±  N02  -J-  NO. 


THE   OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR 


385 


For  the  purpose  of  understanding  its  present  use,  however,  it  is  not 
necessary  to  consider  this  particular  fact.  The  whole  behaves  as  if  it 
contained  nothing  but  N208. 

Details  of  the  Chamber  Process.  —  The  sulphur  dioxide  is 
produced  in  a  row  of  small  furnaces  A  (Fig.  84),  the  structure  of 
which  depends  upon  the  nature  of  the  substance  employed  to  yield 
this  fundamental  constituent  of  sulphuric  acid.  When  good  pyrite 
is  used,  the  ore  burns  unassisted  (p.  42),  while  impure  pyrite  and 
zinc-blende  ZnS  have  to  be  heated,  to  a  greater  or  less  degree,  artifi- 
cially to  maintain  the  combustion.  The  gases  from  the  various  furnaces 
pass  into  one  long  dust  flue,  in  which  they  are  mingled  with  the  proper 


FIG.  84. 


proportion  of  air,  and  have  an  opportunity  to  deposit  oxides  of  iron 
and  of  arsenic  and  other  materials  which  they  transport  mechanically. 
From  this  flue  they  enter  the  Glover  tower  G,  in  which  they  acquire 
the  oxides  of  nitrogen.  Having  secured  all  the  necessary  constituents, 
excepting  water,  and  having  been  reduced  very  considerably  in  tem- 
perature, the  gases  next  enter  the  first  of  the  lead  chambers.  These 
are  large  structures,  from  three  to  five  in  number,  constructed  com- 
pletely of  sheet  lead.  They  vary  in  size,  measuring  as  much  as  100 
X  40  x  40  feet,  and  sometimes  having  a  total  capacity  of  150,000  to 
200,000  cubic  feet.  As  the  gases  drift  through  these  chambers  they 


386  INORGANIC   CHEMISTRY 

are  thoroughly  mixed,  and  an  amount  of  water  considerably  in  excess 
of  that  actually  required,  is  injected  in  the  form  of  steam  at  various 
points.  The  temperature  in  the  first  chamber  is  maintained  at  50°  to 
65°,  while  in  the  last  chamber  it  is  about  15°  above  that  of  the  out- 
side air.  The  acid,  along  with  the  excess  of  water,  condenses  and 
collects  upon  the  floor  of  the  chamber,  while  the  unused  gases,  con- 
sisting chiefly  of  nitrous  anhydride  and  a  very  large  amount  of  nitro- 
gen, derived  from  the  air  originally  admitted,  find  an  exit  into  the 
Gay-Lussac  tower  L. 

This  is  a  tower  about  fifty  feet  in  height,  filled  with  tiles,  over  which 
concentrated  sulphuric  acid  continually  trickles  from  a  reservoir  at 
the  top.  The  object  of  this  tower  is  to  catch  the  nitrous  anhydride 
and  enable  it  to  be  reemployed  in  the  process.  This  is  accomplished 
by  a  reversal  of  action  (2)  above.  The  acid  which  accumulates  in  the 
vessel  at  the  bottom  of  this  tower  contains  the  nitrosylsulphuric  acid, 
and  by  means  of  compressed  air  is  forced  through  a  pipe  up  to  a 
vessel  at  the  top  of  the  Glover  tower  G.  When  this  "  nitrous  vitriol " 
is  mixed  with  dilute  sulphuric  acid  from  a  neighboring  vessel,  by 
allowing  both  to  flow  down  into  the  tower,  the  nitrous  anhydride  is 
once  more  set  free  by  the  interaction  of  the  water  in  the  dilute  acid 
(action  (2)).  The  Glover  tower  is  filled  with  broken  quartz  or  tiles, 
and  the  heated  gases  from  the  furnace  acquire  in  it  their  supply  of 
nitrous  anhydride.  Their  high  temperature  causes  a  considerable 
concentration  of  the  diluted  sulphuric  acid  as  it  trickles  downward. 
The  acid,  after  traversing  this  tower,  is  sufficiently  strong  to  be  used 
once  more  for  the  absorption  of  nitrous  anhydride.  To  replace  the 
nitrous  anhydride  inevitably  lost  by  reduction  to  nitrous  oxide  and 
otherwise,  fresh  nitric  acid  is  furnished  by  small  open  vessels  N, 
containing  sodium  nitrate  and  sulphuric  acid,  placed  in  the  flues  of  the 
py rite-burners.  About  4  kg.  of  the  nitrate  are  consumed  for  every 
100  kg.  of  sulphur. 

The  immense  size  of  the  chambers  is  necessitated  by  the  fact  that 
the  chemical  action,  although  much  quicker  than  the  direct  oxidation 
of  sulphurous  acid,  is  after  all  rather  slow.  The  presence  of  the  large 
amount  of  atmospheric  nitrogen,  which  diminishes  the  concentration  of 
all  the  interacting  substances,  partly  accounts  for  this  slowness.  The 
acid  which  accumulates  upon  the  floors  contains  but  60  to  70  per  cent 
of  sulphuric  acid,  and  has  a  specific  gravity  of  1.5-1.62.  The  excess  of 
water  is  used  to  facilitate  the  second  action.  It  is  required  also  in 
order  that  the  acid  upon  the  floor  may  not  afterwards  absorb  and 


THE   OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR  387 

retain  the  nitrous  anhydride,  for  this  substance  combines  with  an  acid 
containing  more  than  70  per  cent  of  hydrogen  sulphate. 

This  crude  sulphuric  acid  is  applicable  directly  in  some  chemical 
manufactures,  such  as  the  preparation  of  superphosphates  (q.v.).  In 
most  cases,  however,  a  more  concentrated  sulphuric  acid  is  required. 
The  concentration  is  effected  in  the  first  place  by  evaporation  in  pans 
lined  with  lead,  which  are  frequently  placed  over  the  pyrite-burners  in 
order  to  economize  fuel.  The  evaporation  in  lead  is  carried  on  until  a 
specific  gravity  1.7,  corresponding  to  77  per  cent  concentration,  is 
reached.  Up  to  this  point  the  sulphate  of  lead  formed  by  the  action 
of  the  sulphuric  acid  produces  a  crust  which  protects  the  metal  from 
further  action.  The  insoluble  sulphate  of  lead,  however,  becomes 
more  soluble  in  sulphuric  acid  the  more  concentrated  it  is,  and  the 
higher,  therefore,  its  boiling-point.  For  some  applications  of  the  acid, 
concentration  beyond  this  point  is  not  needed.  When  a  stronger  acid 
is  required,  the  water  is  usually  driven  out  by  heating  the  sulphuric 
acid  in  vessels  of  glass  or  platinum.  t  Commercial  sulphuric  acid,  oil  of 
vitriol,  has  a  specific  gravity  1.83-1.84,  and  contains  about  93.5  per 
cent  of  sulphuric  acid. 

In  modern  sulphuric  acid  plants,  the  concentration  beyond  the 
specific  gravity  1.77  is  frequently  carried  out  in  vessels  of  cast  iron,  as 
that  metal  is  attacked  to  an  appreciable  extent  only  by  acids  of 
lower  specific  gravity. 

Physical  Properties.  —  Pure  hydrogen  sulphate  has  a  sp.  gr. 
1.85  at  15°.  When  cooled,  it  crystallizes  (m.-p.  10°).  At  150°-180° 
the  acid  begins  to  fume,  giving  off  sulphur  trioxide.  At  330°  it  boils. 
The  100  per  cent  acid  gives  off  sulphur  trioxide,  and  the  dilute  acid 
water,  until  a  mixture  of  constant  composition  (98.33  per  cent)  and 
constant  b.-p.  330°  is  formed.  This  distils  with  unchanged  compo- 
sition (cf.  p.  182).  The  vapor,  however,  contains  34  per  cent  of  the 
material  in  the  form  of  water  and  sulphur  trioxide,  which  reunite 
when  cooled. 

When  hydrogen  sulphate  is  mixed  with  water  a  considerable  evolu- 
tion of  heat  takes  place.  This  heat  receives  progressively  diminishing 
increments  as  more  water  is  added,  until  a  very  great  dilution  has  been 
reached.  The  total  is  19,400  cal.  This  heat  of  solution  has  not  been 
accounted  for  quantitatively  (cf.  p.  165),  but  a  part  of  it  is  due  to  the 
heat  given  out  in  connection  with  the  ionization  of  the  hydrogen  sulphate. 
The  solution  is  thus  much  more  stable  (i.e.,  it  contains  much  less  energy) 


388  INORGANIC   CHEMISTRY 

than  the  pure  substance,  and  hence  the  latter  absorbs  water  greedily. 
Many  substances  containing  hydrogen  and  oxygen  are  deprived  of 
equivalent  amounts  of  these  elements  by  sulphuric  acid,  just  as  they  are 
by  sulphur  trioxide  (p.  381).  The  same  tendency  is  enlisted  to  promote 
chemical  actions  in  which  water  is  formed,  particularly  in  connection 
with  the  manufacture  of  nitroglycerine  (q.v.),  gun-cotton  (q.v.),  and 
esters  (q.v.). 

Impurities.  —  Commercial  sulphuric  acid  is  frequently  brown  in 
color  on  account  of  the  presence  of  fragments  of  straw  which  have 
become  charred  and  finally  completely  disintegrated.  It  contains  also 
lead  sulphate,  which  appears  as  a  precipitate  when  the  acid  is  diluted, 
as  well  as  arsenic  trioxide  and  oxides  of  nitrogen  in  combination,  and 
many  other  foreign  substances  in  small  quantities.  The  pure  sul- 
phuric acid  employed  in  chemical  laboratories  has  received  special  treat- 
ment for  the  removal  of  these  ingredients. 

• 

Chetnical  Properties  of  Hydrogen  Sulphate*  —  The  compound 
is  not  exceedingly  stable,  for  dissociation  into  water  and  sulphur  triox- 
ide begins  far  below  the  boiling-point  (cf.  p.  387),  and  is  practically 
complete  at  416°,  as  is  shown  by  the  density  of  the  gas.  When 
raised  suddenly  to  a  red  heat  it  is  broken  up  completely  into  water, 
sulphur  dioxide,  and  oxygen. 

When  mixed  with  a  small  quantity  of  water  and  cooled  strongly  it 
gives  a  crystalline  monohydrate  H2S04,  H20,  which  melts  and  is  resolved 
into  its  constituents  at  8°.  Possibly  this  should  be  regarded  as  a  differ- 
ent sulphuric  acid,  H4S06  (cf.  p.  278). 

When  sulphur  trioxide  is  dissolved  in  hydrogen  sulphate,  disul- 
phuric  acid  H2S2O7,  a  solid  compound,  is  obtained.  Hydrogen  sulphate 
containing  80  per  cent  of  disulphuric  acid  is  known  as  "  oleum,"  and  is 
employed  in  chemical  industries.  The  old  "  fuming,"  or  "Nordhau- 
sen  "  sulphuric  acid  contained  10-20  per  cent  of  extra  sulphur  triox- 
ide. The  salts  of  disulphuric  acid  may  be  made  by  strongly  heating 
the  acid  sulphates,  for  example : 

2NaHS04  <=±  Na2S207  +  H20  f . 

In  view  of  this  mode  of  preparation  by  the  aid  of  heat,  they  are  fre- 
quently known  as  pyroaulphates  (Gk.  vvp,  fire).  When  they  are  dig- 
solved  in  water,  the  acid  sulphates  are  reproduced. 

On  account  of  the  large  quantity  of  oxygen  which  hydrogen  sul- 


THE   OXIDES   AND   OXYGEN  ACIDS   OF   SULPHUR  389 

phate  contains,  and  its  instability  when  heated,  it  behaves  as  an  oxi- 
dizing agent.  This  property  has  already  been  illustrated  in  connection 
with  the  action  of  the  acid  upon  carbon  and  sulphur  (p.  379)  and  upon 
copper  (p.  379),  hydrogen  sulphide  (p.  374),  zinc  (p.  96),  and,  particu- 
larly, hydrogen  iodide  (p.  237)  and  hydrogen  bromide  (p.  231).  The 
sulphuric  acid  is  in  consequence  reduced  to  sulphur  dioxide,  and  even 
to  free  sulphur  or  hydrogen  sulphide.  The  inetals,  from  the  most 
active  down  to  silver  (p.  362),  are  capable  of  reducing  it.  Gold  and 
platinum  alone  are  not  attacked,  and  hence  their  use  in  making  sul- 
phuric acid  stills.  Free  hydrogen  itself  is  oxidized  to  water  when 
passed  into  hydrogen  sulphate  at  160° : 

S02(OH)2  +  H2  ->  S02  +  2H2O. 

With  salts  which  it  does  not  oxidize,  hydrogen  sulphate  reacts 
by  double  decomposition  and  sets  free  the  corresponding  acid.  The 
actions  are  always  reversible  ones ;  but  where  the  new  acid  is  volatile, 
as  in  the  case  of  hydrogen  chloride  (p.  180),  we  are  furnished  with 
one  of  the  cheapest  means  of  preparing  acids. 

Since  hydrogen  sulphate  is  dibasic,  that  is,  since  it  has  two  units 
of  hydrogen  which  may  be  replaced  by  a  metal,  it  forms  both  acid 
and  neutral  salts. 

Chemical  Properties  of  Aqueous  Hydrogen  Sulphate.  —  The 

solution  of  sulphuric  acid  is  a  mixture  whose  components  are  :  undisso- 
ciated  molecules  H2SO4,  hydrion  H"7  hydrosulphanion  HS04',  and  sul- 
phanion  SO/'.  The  chemical  properties  shown  by  the  solution  are 
those  of  one  or  other  of  these  components,  according  to  circumstances. 

Except  in  concentrated  solutions  (normal  or  stronger)  the  oxidizing 
effects  of  the  undissociated,  molecular  substance  are  not  encountered. 
The  temperature  of  the  diluted  acid,  even  when  boiling,  is  not  high 
enough  for  the  purpose.  In  fairly  strong  solutions  hydrosulphanion 
is  plentiful  and  shows  itself  in  the  results  of  electrolysis  (see  p.  397). 

The  presence  of  hydrion  is  shown  by  all  its  usual  properties  (p.  345). 
In  the  following  table  the  proportion  of  the  whole  of  the  hydrogen 
existing  in  the  form  of  hydrion  (column  five)  and  its  concentration 
(column  six),  taking  a  normal  solution  of  hydrion  containing  1  g.  per 
liter  as  standard,  are  shown  (cf.  pp.  149  and  329).  The  first  three 
columns  give  the  concentration  of  the  sulphuric  acid  as  a  whole,  in 
terms  (first  column)  of  the  volume  of  liquid  containing  one  equivalent 
( JH2S04  =  49  g.),  in  terms  (second  column)  of  a  normal  solution  as 


390 


INORGANIC   CHEMISTRY 


standard,  and  by  per  cent  (third  column),  respectively.     The  fourth 
column  shows  the  conductivity  (p.  328). 


V 

H2SO< 

PER  CENT 
H2S04. 

A» 

WAoO 

H* 

0.1 

ION 

38.00 

70 

0.18 

1.8N 

1 

N 

4.79 

198 

0.51 

0.51  N 

10 

0.1  N 

0.48 

225 

0.58 

0.058  N 

100 

0.01    N 

0.05 

308 

0.79 

0.0079  N 

1000 

0.001  N 

0.006 

361 

0.93 

0.00093  N 

00 

0 

0.00 

388 

1.00 

0.00 

Column  5  thus  states  that  in  a  normal  solution  51  per  cent,  and  in  a 
centi-normal  solution  79  per  cent  of  the  hydrogen  is  ionic. 

Sulphanion  SO/'  unites  with  all  positive  ions.  The  product,  when 
insoluble,  appears  as  a  precipitate.  The  introduction  of  barium  ions, 
for  example,  by  adding  a  solution  of  barium  nitrate  or  chloride,  is  em- 
ployed as  a  means  of  recognizing  the  presence  of  sulphanion  : 


Naturally  this  test  is  given  by  solutions  of  all  soluble  acid  and  normal 
sulphates.  Since  there  are  other  barium  salts  which  are  insoluble  in 
water,  but  no  common  ones  which  are  not  decomposed  by  acids,  dilute 
nitric  acid  is  first  added  to  the  solution  supposed  to  contain  the  sul- 
phanion. The  other  ions,  if  present,  then  give  no  precipitate  with 
barion. 

Sulphates.  —  The  acid  sulphates,  known  also  as  bisulphates,  may 
be  produced  either  by  semi-neutralization  of  sulphuric  acid  with  a 
base,  followed  by  evaporation  :  NaOH  +  H2S04  <=>  H2O  +  NaHS04,  or 
by  actions  in  which  another  acid  is  displaced,  as  in  making  hydrogen 
chloride  (p.  179). 

The  neutral  (or  normal)  sulphates  are  obtained  by  complete  neutral- 
ization and  evaporation,  or  by  the  second  of  the  above  methods  when 
a  sufficient  amount  of  the  salt  and  a  higher  temperature  are  used  : 

NaCl  +  NaHS04  <=±  Na^SO.+HClf  . 

They  are  often  made  also  by  precipitation,  by  oxidation  of  a  sulphide 
at  a  high  temperature,  PbS  +  2O2  —  »  PbS04,  or  by  addition  of  sulphur 
trioxide  to  the  oxide  of  a  metal  (p.  381). 

Acid  sulphates,  when  heated,  yield  pyrosulphates  (p.  388).  Normal 
sulphates  of  many  heavy  metals  decompose  at  a  white  heat,  giving  off 


THE    OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR  391 

sulphur  trioxide  (p.  380).  The  sulphates  of  potassium,  sodium,  and 
others  of  the  more  active  metals,  however,  are  not  affected  by  heating. 
When  a  sulphate,  or  indeed  any  salt  of  a  sulphur  acid,  is  heated 
strongly  with  carbon,  the  oxygen  is  removed  and  a  sulphide  remains  : 
NajjSO^  +  4C  -»  Na^S  +  4CO.  Upon  this  is  founded  a  general  test  for 
the  presence  of  sulphur  in  any  substance.  The  material  to  be  tested  is 
mixed  with  sodium  carbonate.  A  small  amount  of  the  mixture  is 
placed  on  the  end  of  a  match,  which  has  been  charred  and  rendered 
partially  incombustible  by  previous  application  of  sodium  carbonate. 
When  the  end  of  the  match  is  now  held  in  the  reducing  part  of  the 
Bunsen  flame,  the  compound  of  sulphur,  if  it  contains  oxygen,  is 
reduced  to  the  form  of  sulphide.  This,  by  interaction  with  the  carbon- 
ate, gives  sodium  sulphide,  Na^S.  When  the  product  of  the  reduction 
is  placed  upon  a  silver  coin  and  moistened,  the  sodium  sulphide,  if 
present,  produces  a  black  stain  of  silver  sulphide.  This  is  known  as 
the  hepar  test,  hepar  being  an  old  name  for  a  sulphide. 

Constitution  of  Hydrogen  Sulphate.  —  The  formula  which  we 
assign  to  sulphur  trioxide  is  0  =  S  '     •       It  is  in  general  our  desire 

to  use  the  smallest  possible  valence,  but  here  no  reduction  can  be 
effected  below  the  value  6  for  the  sulphur,  unless  we  join  the  oxygen 

units  to  one  another,  as  in  the  formula  0  =  S  /  V  .   This,  however,  would 

r\  _  TT 

suggest  a  relationship  to  hydrogen  peroxide,   |      -     ,  which  is  not  con- 

0     H 
firmed,  for  hydrogen  peroxide  cannot  be  made  from  sulphuric  acid. 

Assuming,  therefore,  the  above  formula  for  sulphur  trioxide,  the 
addition  of  the  elements  of  water  to  it  in  the  simplest  fashion  results 
in  the  structures  : 

0*0  H-0          O 

^Sf  or  NS^ 

O'l%  H-O'    ^0 

H 

The  second  of  these  two  modes  of  disposing  of  the  water  is  the  one 
which,  in  parallel  cases,  is  usually  most  feasible.  Hardly  any  alter- 
native to  it  is  possible,  for  example,  in  representing  the  action  in 
which  quicklime  is  slaked  (p.  119)  : 


—  (J  ' 


392  INORGANIC   CHEMISTRY 

This  represents  the  change  with  little  derangement  of  the  original 
structure  and  without  alteration  in  the  valence,  while  the  first  unwar- 
rantably increases  the  valence  to  ten.  There  are  other  objections  to 
the  first  formula.  In  it  the  hydrogen  is  supposed  to  unite  more 
immediately  with  the  sulphur,  whereas,  when  the  free  elements  are 
concerned,  hydrogen  actually  combines  more  readily  with  oxygen,  and 
forms  a  more  stable  compound  with  it  than  with  sulphur.  Again, 
compounds  like  hydrogen  sulphide,  H—  S  — H,  in  which  the  hydrogen  is 
undoubtedly  united  to  sulphur,  are  but  slightly  ionized,  and  are  feeble 
acids,  while  hydrogen  sulphate  is  highly  ionized. 

Another  fact  is  more  satisfactorily  accounted  for  by  the  second 
formula.  The  addition  of  chlorine  to  sulphur  dioxide  must  be  shown 
thus: 

.0  Cl  0 

S'     +  €!],->      )g'    , 

^O  Clx    ^O 

for  chlorine  has  a  much  greater  tendency  to  unite  with  sulphur  than 
with  oxygen.  When  the  product,  sulphuryl  chloride,  is  brought  in 
contact  with  water,  sulphuric  acid  and  hydrogen  chloride  are  pro- 
duced. Since  water  has  the  formula  H  — 0  — H,  and  two  molecules  of 
water  are  used,  this  action  is  most  simply  accounted  for,  with  the 
minimum  of  disturbance  in  both  molecules,  by  imagining  the  operation 
to  take  place  as  follows : 


H-0- 
H-O- 


H     Cl  !x 


The  hydrogen  chloride  is  eliminated,  and  the  other  units  of  hydrogen, 
originally  without  doubt  attached  to  oxygen  in  the  water,  may  be  pre- 
sumed to  be  still  connected  with  that  oxygen  when  they  enter  the 
molecule  of  hydrogen  sulphate. 

This  illustration  shows  the  sort  of  reasoning,  based  upon  the 
chemical  properties  and  the  modes  of  formation  of  a  substance,  which 
lead  us  to  the  devising  of  an  appropriate  graphic  or  structural  formula 
(cf.  p.  224).  The  latter  is  not  supposed  in  any  sense  to  represent  the 
actual  physical  structure  of  the  molecule,  but  simply  to  be  a  diagram- 
matic representation  of  the  chemical  relations  of  the  constituents  and 
of  the  chemical  behavior  of  the  whole.  Formulae  of  this  kind  are  in 
continual  use  in  the  study  of  the  compounds  of  carbon,  but  are  seldom 
required  outside  of  that  region. 


THE    OXIDES   AND  OXYGEN   ACIDS   OF   SULPHUR  393 

Uses  of  Sulphuric  Acid.  —  Sulphuric  acid  is  employed  in  almost 
all  chemical  industries,  and  in  some  of  them  (q.v.)  the  quantities 
required  are  very  great.  It  is  used,  for  example,  in  the  first  stage  of 
the  Le  Blanc  process  for  the  manufacture  of  soda,  in  the  refining 
of  petroleum,  in  the  manufacture  of  fertilizers,  in  the  preparation  of 
nitroglycerine  and  gun-cotton,  and  in  the  production  of  coal-tar  dyes. 

OTHER  ACIDS  OP  SULPHUR. 

Hyposulphurous  Acid.  —  The  zinc  salt  of  this  acid  crystallizes 
out  when  zinc  acts  upon  a  solution  of  sulphur  dioxide  in  absolute 
alcohol  : 


Whether  the  acid  is  monobasic  or  dibasic,  and  its  formula  HS02  or 
H2S204,  is  still  uncertain. 

Commercially  a  solution  containing  the  sodium  salt  is  made  by 
the  interaction  of  zinc  with  a  solution  of  sodium  bisulphite  charged 
with  excess  of  sulphur  dioxide  : 


2NaHS03  +  S02  +  Zn  ->  Na,SaO4  +  ZnS03 

The  acid  and  its  salts  are  very  rapidly  oxidized  by  the  air,  the 
former  giving  sulphurous  acid  and  then  sulphuric  acid.  The  above 
solution  of  sodium  hyposulphite  is  used  in  indigo  dyeing,  on  account  of 
its  high  reducing  power.  Indigo,  which  is  insoluble,  is  reduced  by  the 
salt  to  indigo-white  which  passes  into  solution.  When  cloth  saturated 
with  the  mixture,  however,  is  exposed  to  the  air,  the  salt  is  rapidly 
oxidized,  the  indigo-white  likewise  undergoes  oxidation,  and  blue,  in- 
soluble indigo  is  formed  once  more  (see  Dyeing). 

Sulphurous  Acid.  —  This  term  is  applied  to  the  solution  of  sul- 
phur dioxide  in  water.  A  portion  of  the  sulphur  dioxide  remains  dis- 
solved physically,  while  another  portion  is  in  combination  with  the 
water,  forming  sulphurous  acid.  This  in  turn  is  ionized,  and  chiefly, 
after  the  manner  of  the  weaker  dibasic  acids,  into  two  ions,  H"  and 
HS08'.  A  little  S03"  is  formed  from  the  latter.  There  are  thus  in  such 
a  solution  four  mutually  dependent  equilibria  : 

S02(gas)<-»S02(diss'd)  +  H20<=±H2SOS?=>H-  +  HS08'^±H'  +  S08". 

When  the  solution  is  heated,  uncombined  sulphur  dioxide  is  disengaged 
as  a  gas.     The  equilibria  being  thus  disturbed,  the  ions  of  the  acid 


394  INORGANIC   CHEMISTRY 

unite,  the  acid  molecules  decompose,  and  soon  all  the  above  actions  are 
completely  reversed  and  the  whole  of  the  gas  passes  off.  Conversely, 
when  a  base  furnishing  hydroxyl  ions  is  added  to  the  solution  of  the 
acid,  the  hydrogen  ions  disappear,  forming  water,  and  the  above  actions 
all  proceed  in  a  forward  direction  until,  with  a  half -equivalent  of  the 
base,  the  whole  of  the  material  has  been  converted  into  the  form 
HSO8',  in  association,  of  course,  with  the  positive  ions  of  the  base. 
With  a  full  equivalent,  neutralization  follows  and  S03"  is  the  product. 

•  Properties  of  Sulphurous  Acid.  —  The  acid  is  so  unstable  that 
it  cannot  be  obtained  at  ordinary  temperatures  excepting  in  solution 
in  water.  Chemically  it  is  a  comparatively  weak  acid.  As  a  reducing 
agent,  it  is  slowly  oxidized  by  free  oxygen,  and  rapidly  by  oxidizing 
agents,  turning  into  sulphuric  acid.  Thus,  when  free  halogens  are 
added  to  the  solution,  sulphuric  acid  and  the  hydrogen  halide  are  formed : 

H2S03  +  H20  +  I2  «=»  H2S04  +  2HI. 

In  the  particular  case  of  iodine  this  action  takes  place  only  in  very 
dilute  solution,  since  concentrated  sulphuric  acid  decomposes  hydrogen 
iodide  (cf.  p.  237)  and  the  action  is  reversed.  This  interaction  is  used 
in  chemical  analysis  as  a  means  of  estimating  the  quantity  of  sulphur- 
ous acid  in  a  liquid  (cf.  p.  236). 

Hydrogen  peroxide,  potassium  permanganate,  and  other  oxidizing 
agents  convert  the  substance  into  sulphuric  acid  likewise.  It  should 
be  noted  that  in  these  oxidations  we  have,  not  an  addition  of  oxygen  to 
S02,  but  to  the  S08"  or  HS08'  ion  of  the  acid,  whereby  it  passes  into 
the  SO/'  ion  of  sulphuric  acid.  The  ion  is  much  more  easily  oxidized 
than  is  free  sulphur  dioxide  itself. 

Sulphurous  acid  has  the  power  of  uniting  directly  with  many  or- 
ganic coloring  matters  and,  since  the  products  of  this  union  are  usually 
colorless,  it  is  employed  as  a  bleaching  agent.  It  is  especially  useful 
with  materials  like  silk,  wool,  and  straw,  which  are  likely  to  be 
destroyed  by  hypochlorous  acid.  Sunlight  causes  the  dissociation  of 
these  colorless  compounds,  and  so,  with  use,  straw  hats  slowly  recover 
their  original  color.  As  a  disinfectant  it  acts  by  addition  likewise. 

As  a  dibasic  acid,  sulphurous  acid  forms  normal  salts  like  Na2S08, 
and  acid  salts  like  NaHS08. 

Illustration  of  the  Effect  of  Concentration  on  Speed  of  In- 
teraction. —  The  oxidation  of  sulphurous  acid  by  iodic  acid  may  be 
used  to  show  the  effect  of  concentration  on  the  speed  of  an  action 


THE   OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR  395 

(p.  252).  The  iodic  aci<J  may  most  readily  be  made  by  dissolving  po- 
tassium iodate  and  sulphuric  acid  together  in  water,  in  such  quantities 
as  would  give  a  N/2  solution  of  each.  When  1  c.c.  of  this  N/2  iodic 
acid  is  added  to  100  c.c.  of  filtered  starch  emulsion,  and  the  whole  is 
mixed  with  an  equal  volume  of  water  containing  1  c.c.  of  N/2  sulphur- 
ous acid,  the  blue  color  produced  by  the  liberated  iodine  appears  sud- 
denly after  the  lapse  of  a  minute  or  more : 

2HI03  +  5H2S03  ->  SH^SO*  +  H20  +  I2. 

With  double  the  above  quantities  in  the  same  amount  of  water,  that  is, 
with  double  concentrations,  the  speed  of  the  action  is  greatly  increased 
and  the  iodine  becomes  visible  in  less  than  half  the  time. 

Sulphites.  —  The  acid  sulphites  of  the  alkali  metals  (i.e.,  of  potas- 
sium and  sodium)  are  acid  in  reaction,  owing  to  the  appreciable  dissocia- 
tion of  the  ion  HSO/.  The  acid  being  a  weak  one,  however,  solutions 
of  the  normal  salts,  NaoSOg,  etc.,  are  alkaline  towards  litmus  (p.  344). 
The  sulphites  are  readily  decomposed  by  acids  giving  free  sulphurous 
acid,  and  the  latter  partly  decomposes,  yielding  sulphur  dioxide. 

Calcium  bisulphite  solution,  Ca(HS03)2,  is  used  to  dissolve  the  lig- 
nin  out  of  wood  employed  in  the  manufacture  of  paper.  About  30 
per  cent  of  the  wood  is  lignin.  The  rest  is  cellulose  (C6H1006).,.,  and 
constitutes  the  prepared  pulp. 

When  heated,  sulphites  undergo  decomposition.  The  sulphates, 
being  the  most  stable  of  all  the  salts  of  sulphur  acids,  are  formed  when 
the  salts  of  any  of  those  acids  are  decomposed  by  heating.  The  nature 
of  the  particular  salt  determines  what  other  products  shall  appear. 
Here,  one  molecule  of  the  sulphite  furnishes  three  atoms  of  oxygen, 
sufficient  to  oxidize  three  other  molecules,  and  leaves  one  molecule  of 
sodium  sulphide  behind : 

4Na2S08  ->  Na2S  +  3Na*SO4. 

The  acid  sulphites  (bisulphites)  first  lose  sulphurous  acid,  before  chang- 
ing in  this  way.  Thus,  sodium  hydrogen  sulphite  begins  by  decompos- 
ing as  follows : 

2NaHS08  -»  Na,,S08  +  H2SOS  (or  H2O  +  S02). 

The  acid  salts  of  volatile  acids,  when  heated,  all  decompose  in  this 

(tf-  PP-  241,  388). 
The  sulphites  are  as  readily  oxidized  as  is  the  acid  itself.    They  are 


396  INORGANIC   CHEMISTRY 

slowly  converted,  both  in  solution  and  in  the  solid  form,  by  the  influ- 
ence of  the  oxygen  of  the  air,  into  sulphates.  It  is  interesting  to  note 
that  the  addition  of  sugar  or  glycerine  to  a  solution  of  a  sulphite  re- 
duces the  speed  of  oxidation  by  free  oxygen  very  markedly.  These 
substances  act  as  catalytic  agents  ;  and  the  present  case  shows  that 
agents  of  this  kind  may  not  only  increase  the  speed  of  actions,  which 
is  their  usual  function,  but  may  also  have  a  restraining  influence. 

Thio  sulphuric  Acid.  —  This  acid  is  not  known  in  the  free  con- 
dition, but  its  salts  are  in  common  use  in  the  laboratory  and  commer- 
cially. The  sodium  salt,  for  example,  is  prepared  by  boiling  a  solu- 
tion of  sodium  sulphite  with  free  sulphur.  The  action  is  something 
like  the  addition  of  oxygen  to  sulphurous  acid  : 

NaJSO,  +  S  ->  NaAO,     or     SO."  +  S  -»  8,0,"- 

The  product,  thiosulphate  of  sodium,  is  used  in  photography  as  a 
solvent  for  salts  of  silver,  and  is  commonly  known  as  hyposulphite  of 
soda.  Its  solution  in  water  forms  the  fixing  bath  of  the  photographer. 
By  the  addition  of  acids  to  a  solution  of  sodium  thiosulphate,  the 
thiosulphuric  acid  is  set  free,  but  the  latter  instantly  decomposes  : 

,  4-  2HC1  «=»  H2S208  +  2NaCl, 


Even  carbon  dioxide  from  the  air,  giving  carbonic  acid  (q*v.\  pro- 
duces this  effect  slowly  in  fixing  solutions.  The  actions  being  rever- 
sible, preliminary  addition  of  a  sulphite  to  the  solution  helps  to  sustain 
the  reverse  action,  in  which  sulphurous  acid  is  a  factor,  and  so  pre- 
serves the  solution.  The  delay  in  the  appearance  of  the  precipitate 
of  sulphur  in  dilute  solutions  is  due  to  the  temporary  existence  of  a 
supersaturated  solution  (cf.  p.  160)  of  the  free  element. 

Iodine  acts  upon  sodium  thiosulphate  solution,  giving  sodium 
tetrathionate  : 

2Na,Sa08  +  I2  -»  2NaI  +  Na2S406. 

This  action  is  used,  by  employment  of  a  standard  solution  of  sodium 
thiosulphate,  for  estimating  quantities  of  free  iodine  in  analysis.  The 
disappearance  of  the  color  of  the  latter  indicates  that  a  sufficient 
amount  of  the  salt  has  been  employed.  When  chlorine  is  used,  the 
oxidation  is  more  complete.  The  products  are  sodium  sulphate,  sul- 
phuric acid,  and  hydrochloric  acid  : 

,  +  5H20  +  4CL,  ->  N^SO^-h  H2S04  +  8HC1. 


THE   OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR  397 

In  consequence  of  the  very  great  amount  of  free  chlorine  which  the 
sodium  thiosulphate  is  thus  able  to  transform,  it  is  employed,  as 
j  for  the  purpose  of  removing  chlorine  from  bleached  fabrics. 


Persulphuric  Acid.  —  This,  like  the  other  acids  just  mentioned, 
is  unstable,  and  can  be  kept  only  in  very  dilute  solution.  Its  salts, 
however,  are  coming  into  use  for  commercial  purposes  and  for  "  redu- 
cing "  negatives  in  photography.  When  a  discharge  of  electricity  is 
passed  through  a  mixture  of  sulphur  trioxide  and  oxygen,  drops  of 
liquid  are  formed  which  appear  to  have  the  composition  S207,  and 
when  dissolved  in  water  give  dilute  persulphuric  acid,  S207  -f-  H20  —  > 
H2S208.  More  significant  of  its  relations  is  its  formation,  to  some 
extent,  when  concentrated  sulphuric  acid  and  a  strong  solution  of 
hydrogen  peroxide  are  mixed  : 

2H2S04  -f  HA  <=>  H2S208  +  2H.O. 

This  action  is  reversible.*  Under  some  circumstances,  monopersul- 
phuric  acid  is  formed:  H2S04-f-H202<=>  H2SO5-|-H20.  Interesting  in 
its  way,  also,  is  its  production  in  the  electrolysis  of  aqueous  sulphuric 
acid: 

2HS04  +  2  0  -»  H2S208. 

This  action  is  most  conspicuous  in  rather  concentrated  solutions  in 
which  hydrosulphanion  is  plentiful  (cf.  p.  389)  and  when  a  small  anode, 
resulting  in  severe  crowding  of  the  HS04  radicals  as  they  are  liber- 
ated, is  employed.  The  salts  are  prepared  by  electrolyzing  sodium 
hydrogen  sulphate  NaHS04  in  concentrated  solution.  The  persul- 
phuric acid,  formed  by  the  union  of  the  negative  ions  in  pairs,  under- 
goes double  decomposition  with  the  excess  of  sodium  bisulphate,  and 
the  less  soluble  sodium  persulphate  crystallizes  out.  The  other  salts 
are  made  by  double  decomposition  from  this  one. 

*  This  action  and  the  next  are  not  easily  classifiable  under  any  of  the  ten 
kinds  formerly  discussed  (p.  187).  They  consist  in  the  union  of  H  and  OH  to  form 
water  : 


HS04:  H  +  HO  ;  OH  +  H  |SO4H  ->  2H.P  +  (S04H)2. 

Neutralizations  (p.  851)  they  are  not,  because  the  interacting  substances  are  both 
acids.  Just  as  the  loss  of  water  from  one  acid  gives  an  anhydride,  so  here,  the 
loss  of  water  between  two  acids  gives  a  mixed  anhydride  (see  Cblorosulphuric 
acid,  below). 


398  INORGANIC   CHEMISTRY 

The  persulphates  decompose  readily  when  heated,  yielding  pyro. 
sulphates  and  oxygen : 


The  solution  of  the  acid  is  an  active  oxidizing  agent : 
H2S208  +  H20  •++  2H2S04  ( +  0). 

Polythionic  Acids.  —  Di-,  tri-,  tetra-,  and  pentathionic  acid  (p. 
381)  are  all  formed  simultaneously  when  sulphur  dioxide  and  hydro- 
gen sulphide  gases  are  passed  alternately  into  water,  although  the 
gases  themselves  (p.  373)  interact  to  produce  simply  free  sulphur  and 
water : 

H2S-f-3S02->H2S406, 
2H2S  +  6S02  -»  H2S306  +  H2S606, 
3H2S  +  9S02  _>  H2S206  +  2H2S5O6. 

Most  of  these  acids  and  their  salts  are  of  minor  interest  and  need  not 
be  discussed. 

The  production  of  sodium  tetrathionate  by  the  action  of  iodine 
upon  sodium  thiosulphate  has  already  been  mentioned. 

When  manganese  dioxide  is  treated  with  sulphurous  acid,  it  inter- 
acts very  rapidly  and  a  solution  of  manganous  dithionate  is  obtained : 

Mn02  +  2H2S08  -»  MnS206  +  2H2O. 

The  salts  of  these  acids  are  in  many  cases  fairly  stable,  but  the 
acids  themselves  decompose  readily  when  set  free. 

COMPOUNDS  OF  SULPHUR  AND  CHLORINE. 

Sulphur  Monochloride.  —  When  chlorine  gas  is  passed  over 
heated  sulphur  it  is  absorbed,  and  a  dark  reddish-yellow  liquid,  boiling 
at  138°,  is  obtained.  The  molecular  weight  of  this  substance,  as  shown 
by  the  density  of  its  vapor,  indicates  that  it  possesses  the  formula 
S8C12.  When  thrown  into  water  it  is  rapidly  hydrolyzed,  producing 
sulphur  dioxide  and  free  sulphur  : 

2S2C12  +  2H20  -»  S02  +  4HC1  +  38. 

Sulphur  itself  dissolves  very  freely  in  the  monochloride,  and  the 
solution  is  employed  in  vulcanizing  rubber. 


THE   OXIDES   AND   OXYGEN   ACIDS   OF   SULPHUR  399 

By  surrounding  sulphur  monochloride  with  a  freezing  mixture, 
and  treating  it  with  excess  of  chlorine,  a  liquid  tetrachloride  SC14  is 
formed.  This  gives  up  chlorine,  and  turns  into  the  monochloride 
again,  when  allowed  to  become  warm. 

Thionyl  Chloride.  —  By  the  action  of  sulphur  dioxide  gas  upon 
phosphorus  pentachloride,  part  of  the  oxygen  in  the  former  is  replaced 
by  chlorine  : 

S03  +  PC15  !-*  SOCL,  +  POC1,. 

The  products  are  thionyl  chloride  and  phosphorus  oxychloride.  The 
former  is  a  colorless  liquid,  boiling  at  78°,  and  is  separated  from  the 
latter  (b.-p.  107°)  by  fractional  distillation  (see  Petroleum).  It  is  de- 
composed immediately  on  contact  with  water  : 

SOC12  +  H20  ->  S02  +  2HC1. 

Sulphuryl  Chloride.  —  Sulphur  dioxide  and  chlorine  gases  unite 
when  exposed  to  direct  sunlight  to  form  a  liquid  known  as  sulphuryl 
chloride  S02C12.  When  camphor  is  introduced  into  the  vessel  the 
union  takes  place  much  more  rapidly,  owing  to  some  catalytic  effect 
of  this  substance.  The  compound  is  a  colorless  liquid,  boiling  at  69°. 
With  water  it  gives  sulphuric  acid  and  hydrogen  chloride  (p.  392). 
When  a  strictly  limited  amount  of  water  is  supplied,  a  partial  action 
of  the  same  nature  occurs,  and  the  product  is  known  as  chlorosulphuric 
acid: 

80,01,  +  H30  -»  S0a  +  HC1. 


This  intermediate  compound  may  be  formed  also  by  the  addition  of 
hydrogen  chloride  to  sulphur  trioxide. 

Exercises.  —  1.   What  ground  is  there  for  assigning  the  formula 
S02  instead  of  S204  to  sulphur  dioxide  (p.  193)  ? 

2.  Explain  why  nitric  acid  is  completely  displaced  by  the  action  of 
sulphuric  acid  on  sodium  nitrate  (p.  384). 

3.  How  many  times,   on  an  average,  does  a  molecule  of  nitrous 
anhydride  go  through  the  cycle  of  changes  by  which  sulphuric  acid  is 
produced  before  it  is  eliminated  in  some  other  form  (p.  386)  ? 

4.  Make  a  list  of,  and  classify,  the  various  applications  of  sulphuric 
acid  to  the  liberation  of  other  acids. 


400  INORGANIC   CHEMISTRY 

5.  Formulate  the  behavior  of  the  hydrosulphanion  (p.  389)  when 
a  solution  of  barium  chloride  is  added  to  a  rather  concentrated  solu- 
tion of  sulphuric  acid. 

6.  Can  you  give  any  reasons  for  preferring  to  regard  KHS04,  and 
substances  like  it,  as  acid  salts  rather  than  double  salts  of  the  form 
K2SO4,  H2SO4  (p.  358)  ? 

7.  From  a  consideration  of  the  physical  conditions,  how  can  you 
account  for  the  fact  that  hyposulphurous  acid  is  oxidized  by  free 
oxygen  first  to  sulphurous  acid  and  then  to  sulphuric  acid  (p.  393), 
while  moist  free  sulphur,  even  with  oxidizing  agents,  gives  sulphuric 
acid  directly  (p.  371)  ? 

8.  Why  were  not   disulphuric  acid   and   monopersulphuric   acid 
placed  in  the  list  on  p.  381  (cf.  p.  278)  ? 

9.  Write  in  ionic  form  the  equation  for  the  interaction  of  sodium 
thiosulphate  and  iodine  in  aqueous  solution. 

10.  Restate  the  import  of  the  following  sentence  in  such  a  way  as 
to  avoid  the  use  of  the  hypothesis  of  ions  :  Liquefied  sulphur  dioxide 
"  ionizes  substances  dissolved  in  it  as  well  as  does  water  "  (p.  379). 


CHAPTER    XXIII 
SELENIUM    AND    TELLURIUM  :    THE    PERIODIC    SYSTEM 

ALONG  with  sulphur,  chemists  group  two  other  elements,  selenium 
(Se,  at.  wt.  79.2)  and  tellurium  (Te,  at.  wt.  127.6).  If  the  nature  of 
the  chief  compounds  of  sulphur  is  kept  in  mind,  the  close  analogy 
between  the  nature  and  chemical  behavior  of  the  three  elements  and 
their  corresponding  compounds  will  be  noticed  at  once  (see  Chemical 
relations  of  the  sulphur  family,  below). 

SELENIUM. 

Occurrence  and  Properties  of  the  Element.  —  Selenium  (Gk. 
o-eA^i/Tj,  the  moon)  occurs  free  in  some  specimens  of  native  sulphur, 
and  in  combination  often  takes  the  place  of  a  small  part  of  the  sulphur 
in  pyrite  (FeS2).  It  is  found  free  in  the  dust-flues  of  the  pyrite- 
burners  of  sulphuric  acid  works.  The  familiar  forms  are,  the  red  pre- 
cipitated variety,  which  is  amorphous  and  soluble  in  carbon  disulphide, 
and  the  lead-gray,  semi-metallic  variety,  obtained  by  slow  cooling 
of  melted  selenium,  which  is  insoluble,  and  melts  at  217°.  In  the 
latter  form  it  has  some  capacity  for  conducting  electricity,  which  is 
increased  by  exposure  to  light  in  proportion  to  the  intensity  of  the 
illumination.  It  boils  at  680°,  and  at  high  temperatures  has  a  vapor 
density  corresponding  to  the  formula  Se2. 

The  element  unites  directly  with  many  metals,  burns  in  oxygen  to 
form  selenium  dioxide,  and  unites  vigorously  with  chlorine. 

Hydrogen  Selenide.  —  Ferrous  selenide,  made  by  heating  iron 
filings  with  selenium,  when  treated  with  concentrated  hydrochloric 
acid  gives  hydrogen  selenide  : 

FeSe  +  2HC1  <±  H^ef  +  FeCL,. 


The  compound  is  a  poisonous  gas,  which  possesses  an  odor  recalling 
rotten  horse-radish,  and  is  soluble  in  water.  The  solution  is  faintly 
acid  in  reaction,  and  deposits  selenium  when  exposed  to  the  action  of 

401 


402  INORGANIC   CHEMISTRY 

the  air  (cf.  p.  373).  Other  selenides,  which,  with  the  exception  of 
those  of  potassium  and  sodium,  are  insoluble  in  water,  may  be  pre- 
cipitated by  leading  the  gas  into  solutions  of  soluble  salts  of  appro- 
priate metals  (cf.  p.  375). 

Selenium  Dioxide  and  Selenious  Acid.  —  The  dioxide  (Se02) 
is  a  solid  body  formed  by  burning  selenium  or  evaporating  a  solution 
of  selenious  acid  H2Se08.  The  latter  may  be  made  directly  by  oxidiz- 
ing selenium  with  boiling  nitric  acid  or  aqua  regia  (q-v.).  Unlike 
sulphur  (p.  371),  the  element  gives  little  of  the  higher  acid  H2Se04  by 
this  treatment.  In  aqueous  solution  the  acid  is  easily  reduced  to 
selenium.  In  fact,  sulphurous  acid  can  be  oxidized  by  it  : 


H2Se08  +  211,80,  ->  211,804  +  H2O  +  Se. 

Selenic  Acid.  —  No  trioxide  is  known.  Selenic  acid  (H2Se04)  is 
made  by  using  the  most  powerful  oxidizing  agents  with  an  aqueous 
solution  of  selenious  acid.  It  can  be  isolated  as  a  white  solid.  It  is 
itself  a  powerful  oxidizing  agent,  and,  even  in  dilute  solution,  liberates 
chlorine  from  hydrochloric  acid  : 

H2Se04  +  2HC1  -4  ILjSeOg  +  H20  +  CL,. 

Sulphuric  acid  (cf.  p.  389),  on  the  other  hand,  is  an  oxidizing  agent 
only  in  somewhat  concentrated  form,  and  even  then  it  can  oxidize 
hydrobromic  acid  (p.  231),  but  not  hydrochloric  acid. 

Chlorides  of  Selenium.  —  The  chlorides  are  formed  by  direct 
union  of  the  elements.  The  tetrachloride,  a  yellow  crystalline  sub- 
stance, is  formed  when  the  monochloride  is  heated  :  2Se2Cl2  —  •»  3Se  4- 
SeCl4,  the  behavior  in  the  case  of  sulphur  chlorides  being  just  the 
inverse  of  this  (p.  399). 

TELLURIUM. 

Tellurium  (Lat.  tellus,  the  earth)  occurs  in  sylvanite  in  combina- 
tion with  gold  and  silver.  It  is  a  white,  metallic,  crystalline  sub- 
stance, melting  at  452°.  When  formed  by  precipitation  it  is  a  black 
powder.  It  conducts  electricity  to  some  extent.  The  vapor  density 
corresponds  to  the  formula  Te2. 

The  free  element  unites  with  metals  directly,  and  burns  in  air  to 
form  the  dioxide. 


SELENIUM   AND   TELLURIUM:   THE   PERIODIC   SYSTEM       403 

Hydrogen  telluride  (H2Te)  is  made  by  the  action  of  acids  on 
metallic  tellurides,  and  its  aqueous  solution  is  rapidly  oxidized  by  air 
with  precipitation  of  tellurium.  The  tellurides  of  the  alkali  metals 
are  soluble  in  water,  the  others  are  insoluble. 

Tellurious  acid  (H2Te08)  is  formed  by  oxidizing  the  element  with 
nitric  acid,  and  is  a  crystalline  solid,  little  soluble  in  water.  It  is  a 
feeble  acid,  of  which  many  salts  have  been  made.  It  is  also  some- 
what basic,  a  sulphate  (2Te02,  S08)  and  a  nitrate  (Te208(OH)ISr08)  being 
known.  In  this  respect  it  differs  markedly  from  sulphurous  acid. 

Telluric  acid  is  made  by  oxidizing  tellurious  acid  in  aqueous  solu- 
tion with  chromic  acid  (p.  374).  It  is  difficultly  soluble  in  water.  It 
does  not  affect  indicators,  and  is  therefore  actually  more  feebly  acidic 
than  is  hydrogen  sulphide.  It  is  obtained  as  a  solid  having  the  com- 
position H6Te06  (or  3H20,  Te08)  on  evaporating  the  solution.  When 
heated,  this  body  loses  water,  some  of  the  trioxide  (Te08)  being 
formed.  The  last  is  a  yellow  solid,  which  shows  no  tendency  to 
recombine  with  water,  in  this  respect  resembling  silica  (q.v.).  Tel- 
lurates  of  the  alkali  metals  may  be  made  by  heating  the  tellurites 
with  potassium  or  sodium  nitrate :  K2Te08  -f-  KN08  — >  K^TeC^  -|- 
KN02. 

Tellurium  forms  two  very  stable  chlorides,  TeCIa  and  TeCl4,  which 
are  decomposed  by  water.  The  second,  however,  exists  in  solution 
with  excess  of  hydrogen  chloride  :  TeCl4+3H2O  ^±  H2Te08  +  4HC1, 
showing  the  tellurious  acid  to  be  basic  in  properties  and  the  element 
tellurium  to  be,  to  a  certain  degree,  a  metal  (see  Chap,  xxxii). 

The  Chemical  Relations  of  the  Sulphur  Family. — It  will  be 
seen  that  sulphur,  selenium,  and  tellurium  are  bivalent  elements  when 
combined  with  hydrogen  or  metals.  In  combination  with  oxygen  they 
form  unsaturated  compounds  of  the  form  XIV02,  while  their  highest 
valence  is  found  in  S08,  Te03,  and  H2Se04,  where  they  must  be  sexi- 
valent.  The  general  behavior  of  corresponding  compounds  is  very 
similar.  At  the  same  time,  there  is  in  all  cases  a  progressive  change  as 
we  proceed  from  sulphur  through  selenium  to  tellurium.  The  elemen- 
tary substances  themselves,  for  example,  become  more  like  metals, 
physically,  and  they  show  higher  and  higher  melting-points.  The 
affinity  for  hydrogen  decreases,  as  is  shown  by  the  increasing  ease 
with  which  the  compounds  H^  are  oxidized  in  air.  The  affinity  for 
oxygen  likewise  decreases,  for  the  elements  become  increasingly  diffi- 
cult to  raise  to  the  highest  state  of  oxidation.  On  the  other  hand, 


404  INORGANIC   CHEMISTRY 

the  tendency  to  form  higher  chlorides  becomes  greater.  We  note  also 
that  the  compounds  H2XO4  become  less  and  less  active  as  acids,  and  a 
basic  tendency  begins  to  assert  itself. 

THE  PERIODIC  SYSTEM. 

Classification,  or  the  arrangement  of  facts  on  the  basis  of  likeness, 
is  part  of  the  method  of  science.  In  chemistry  the  multitude  of  facts 
is  not  less  than  in  other  sciences,  and  the  necessity  of  arrangement 
equally  urgent.  It  is  needed  to  make  possible  the  systematic  descrip- 
tion of  the  ascertained  facts,  and  to  furnish  a  guide  in  investigation,  by 
suggesting  stochastic  hypotheses,  and  so  pointing  out  directions  in 
which  new  facts  of  interest  may  be  found.  Thus,  we  have  treated  the 
halogens  as  a  group ;  and  chemists,  knowing  how  hypochlorous  acid 
(HC10)  and  perchloric  acid  (HC104),  and  their  salts,  are  made,  have 
been  led  to  attempt  to  obtain  related  substances,  like  HIO  and  HBr04 
and  their  salts,  by  methods  suggested  by  analogy. 

At  first  sight,  the  most  definite  method  of  classification  would 
appear  to  be  the  grouping  of  elements  of  like  valence.  But  this  brings 
together  sodium  and  chlorine  —  an  element  whose  hydrogen  compound 
is  unstable  and  without  markedly  characteristic  properties,  and  whose 
hydroxyl  compound  is  an  active  base,  with  an  element  whose  hydrogen 
compound  is  an  active  acid  and  whose  hydroxyl  compound  is,  in  a 
feeble  degree,  an  acid  also.  This  method  homologates  similar  and  con-* 
trasting  elements  indiscriminately. 

Metallic  and  Non-Metallic  Elements.  —  Thus  far  we  have  found 
the  division  into  metallic  and  non-metallic  elements  very  serviceable  for 
classification  in  terms  of  chemical  relations  (p.  177).  This  distinction 
we  shall  continue  to  employ.  The  metals,  or  positive  elements  (p.  119), 
(1)  form  positive  ions  containing  no  other  element  (cf.  p.  337).  Thus 
the  metals  give  sulphates,  nitrates,  carbonates,  and  other  salts,  which 
furnish  a  metallic  ion  together  with  the  ions  SO/',  N03',  and  CO8".  (2) 
Their  hydroxides,  KOH,  Ca(OH)2,  etc.,  give  the  same  metallic  ion,  and 
the  rest  of  the  molecule  forms  hydroxidion.  That  is  to  say,  their 
hydroxides  are  bases  and  their  oxides  are  basic.  They  often  enter 
with  other  elements  into  the  composition  of  a  negative  ion,  as  is  the  case 
with  manganese  in  K.Mn04,  with  chromium  in  K2.O207,  and  with  silver 
in  K.Ag(CN)2.*  But  the  most  definitely  metallic  elements  form  with 

*  The  mode  of  division  into  ions  is  shown  by  the  position  of  the  period  in  the 
formula. 


SELENIUM   AND   TELLURIUM:   THE   PERIODIC   SYSTEM       405 

oxygen  such  a  negative  ion  only  while  exhibiting  a  different  valence 
from  that  which  they  possess  when  acting  as  positive  elements.  Thus, 
manganese  when  a  positive  element  has  the  valences  two  and  three, 
MnO  and  Mn203,  Mn.Cl2  and  Mn.Cl8,  Mn.S04  and  Mn2.(S04)8,  etc., 
while  in  permanganates  we  have  potentially  the  oxide  Mn2O7,  (K20, 
Mn2O7  s=  2KMnO4),  in  which  it  is  heptavalent.  The  graphic  formulae 
express  the  difference  in  valence  : 

ri  xcl  ^° 

Mn/  Mn-Cl  K-O-Mn=O 


Manganese  forms  no  heptachloride  or  heptanitrate  (corresponding  to 
the  heptoxide)  in  which  manganese  alone  would  form  a  heptavalent 
positive  ion  on  the  one  hand,  and  no  compound  in  which  bivalent  or 
trivalent  manganese  with  oxygen  form  the  negative  ion,  on  the  other. 
Chromium  (q.v.)  is  bivalent  or  trivalent  when  a  metal,  and  sexivalent 
when  acting  as  a  non-metal.  The  various  hydroxides  of  these  elements 
are  bases  or  acids  in  accordance  with  this  distinction.  If  we  knew  only 
the  compounds  in  which  manganese  and  chromium  are  heptavalent  and 
sexivalent,  respectively,  we  should  regard  them  as  non-metallic  ele- 
ments pure  and  simple,  the  metallic  appearance  of  the  free  elements  to 
the  contrary  notwithstanding. 

The  non-metals  or  negative  elements,  (1)  are  found  chiefly  in  nega- 
tive ions.  They  form  no  nitrates,  sulphates,  carbonates,  etc.,  for  they 
could  not  do  so  without  themselves  constituting  the  positive  ion.  We 
have  no  such  salts  of  oxygen,  sulphur,  carbon,  or  phosphorus,  for  ex- 
ample. (2)  Their  hydroxides,  like  C102OH,  P(OH)3,  S02(OH)2,  furnish 
no  hydroxyl  ions,  as  this  would  involve  the  same  consequence.  These 
hydroxides  are  divided  by  dissociation,  in  fact,  so  that  the  non-metal 
forms  part  of  a  compound  negative  radical,  and  the  other  ion  is  hydrion, 
C103.H,  P03H.H2,  S04.H2.  (3)  Their  halogen  compounds,  like  PCI,, 
(p.  181)  and  S^lg  (p.  398),  are  completely  hydrolyzed  by  water,  and  the 
actions  are  not,  in  general,  reversible.  The  halides  of  the  typical 
metals  are  not  hydrolyzed  (see  Chap,  xxxii) 

The  distinction  is  not  perfectly  sharp,  however.  Thus,  zinc  (q.v.) 
gives,  both  salts  like  the  sulphate,  Zn.S04,  and  chloride,  Zn.CLy  and  com- 
pounds like  sodium  zincate  (p.  99),  ZnO^Na-j,  showing  the  same  valence 
in  both  classes  : 

rj    /Cl  -    /0-Na 

Zn\Cl  Zn  \0-Na 


406  INORGANIC   CHEMISTRY 

Its  hydroxide  ionizes  in  two  ways,  Zn.(OH)2  and  ZnO^K,.  Similarly 
tellurious  acid,  H2Te08,  acts  both  as  acid  and  base  (p.  403).  We  shall 
find  this  double  behavior  conspicuous  in  the  compounds  of  arsenic  and 
antimony.  In  spite  of  the  partial  merging  of  the  two  classes  of  ele- 
ments, however,  the  general  distinction  is  worth  preserving  (see  Chap, 
xxxii). 

Classification  by  Atomic  Weights.  —  A  closer  discrimination 
than  that  furnished  by  these  two  categories  is  required,  however,  and 
a  study  of  the  order  into  which  the  elements  fall  when  arranged  accord- 
ing to  their  atomic  weights  has  provided  this. 

The  first  indication  of  a  significant  relation  between  the  atomic 
weights  and  the  properties  of  the  elements  was  given  by  a  fact  noted 
by  Dobereiner  (1829).  He  drew  attention  to  the  existence  of  closely 
similar  elements  in  sets  of  three  (triads),  where  the  central  element 
was  intermediate  in  properties  between  the  two  others,  and  the  atomic 
weight  of  the  central  element  was  almost  the  exact  arithmetical  mean 
of  the  weights  of  the  other  two.  The  three  following  triads  illustrate 
this  relation : 

Chlorine    .     .     .     35.45  Sulphur      .     .     .     32.06  Calcium  ....     40.1 

Bromine    .     .     .     79.96  Selenium    .     .     .     79.2  Strontium     .     .     .    87.6 

Iodine        .     .     .  126.97  Tellurium   .     .     .  127.6  Barium    ....  137.4 

Mean  of  Cl  and  I,     81.1  Mean  of  S  and  Te,   79.8  Mean  of  Ca  and  Ba,   88.7 

Newlands  (1863-4)  called  attention  to  the  existence  of  a  surprising 
regularity  when  the  elements  were  placed  in  the  order  of  ascending 
atomic  weight.  Omitting  hydrogen  (1)  the  first  seven  are  :  lithium  (7), 
glucinum  (9),  boron  (11),  carbon  (12),  nitrogen  (14),  oxygen  (16), 
fluorine  (19).  These  are  all  of  totally  different  classes,  and  include  first 
a  metal  forming  a  strongly  basic  hydroxide,  then  a  metal  of  the  less 
active  sort,  then  five  non-metals  of  increasingly  negative  character,  the 
last  being  the  most  active  non-metal  known.  The  next  element  after 
fluorine  (19)  is  sodium  (23),  which  brings  us  back  sharply  to  the 
elements  that  form  strongly  basic  hydroxides.  Omitting  none,  the  next 
seven  elements  are :  sodium  (23),  magnesium  (24.4),  aluminium  (27), 
silicon  (28.4),  phosphorus  (31),  sulphur  (32),  chlorine  (35.5). 

In  this  series  there  are  three  metals  of  diminishing  positiveness, 
followed  by  four  non-metals  of  increasing  negative  activity,  the  last 
being  a  halogen  very  like  fluorine.  On  account  of  the  fact  that  each 
element  resembles  most  closely  the  eighth  element  beyond  or  before  it 
in  the  list,  the  relation  was  called  the  law  of  octaves.  After  chlorine 


SELENIUM   AND   TELLURIUM:   THE   PERIODIC   SYSTEM      407 

the  octaves  become  less  easy  to  trace.  Potassium  (39)  follows  chlo- 
rine and  corresponds  satisfactorily  to  sodium,  but  it  is  not  until  seven- 
teen successive  elements  have  been  set  down  that  we  reach  one  closely 
resembling  chlorine,  namely,  bromine. 

That  this  periodicity  in  chemical  nature  is  more  than  a  coincidence 
is  shown  by  the  fact  that  the  valence  and  even  the  physical  properties, 
such  as  the  specific  gravity,  show  a  similar  fluctuation  in  each  series, 
and  recurrence  in  the  following  one.  In  the  first  two  series  the  com- 
pounds with  other  elements  are  of  the  types  : 

LiCl,   010.,   BC1.,  CO., 


Thus  the  valence  towards  chlorine  or  hydrogen  ascends  to  four  and 
then  reverts  to  one  in  each  octave.  The  highest  valence,  shown  in 
oxygen  compounds,  ascends  from  lithium  to  nitrogen  with  values  one 
to  five,  and  then  fails  because  compounds  are  lacking.  In  the  second 
octave,  however,  it  goes  up  continuously  from  one  to  seven. 

Again,  the  specific  gravities  of  the  elements  in  the  second  series, 
using  the  data  for  red  phosphorus  and  liquid  chlorine,  are  : 

Na  0.97,  Mg  1.75,  Al  2.67,   Si  2.49,  P  2.14,   S  2.06,  01  1.33. 

Of  greater  significance  chemically  are  the  related  numbers  represent- 
ing the  volumes  in  cubic  centimeters  occupied  by  a  gram-atomic  weight 
of  each  element  (the  atomic  volumes)  : 

Na  24,  Mg  14,  Al  10,   Si  11,  P  14,   S  16,   Cl  27. 

A  similar  regular  fluctuation  is  shown  by  all  the  physical  properties  of 
corresponding  compounds. 

Mendelejeff's  Scheme.  —  In  1869  Mendelejeff  published  an  im- 
portant contribution  towards  adjusting  the  difficulty  which  the  ele- 
ments following  chlorine  presented,  and  developed  the  whole  concep- 
tion so  completely  that  the  resulting  system  of  classification  has  been 
connected  with  his  name  ever  since.  Almost  simultaneously  Lothar 
Meyer  made  similar  suggestions,  but  did  not  urge  them  with  the  same 
conviction  or  elaborate  them  so  fully.  The  following  table,  in  which 
the  atomic  weights  are  expressed  in  round  numbers,  is  a  modification 
of  one  of  Mendele  Jeff's. 


408 


INORGANIC   CHEMISTRY 


Ng 


'28 


Ss 


oS 


2 


$ 


*- 


SELENIUM    AND   TELLURIUM:    THE   PERIODIC   SYSTEM      409 

The  chief  change  from  the  arrangement  in  simple  octaves  is  that  the 
third  series,  beginning  with  potassium,  is  made  to  furnish  material  for 
two  octaves,  potassium  to  manganese  and  copper  to  bromine,  and  is 
called  a  long  series.  The  valences  fall  in  with  this  plan  fairly  well. 
Copper,  while  usually  bivalent,  forms  also  a  series  of  compounds  in 
which  it  appears  to  be  univalent.  Iron,  cobalt,  and  nickel  cannot  be 
accommodated  in  either  octave,  as  their  valences  are  always  two  or 
three.  At  the  time  Mendelejeff  made  the  table,  three  places  in  the 
third  series  had  to  be  left  blank,  as  a  trivalent  element  [Sc]  was  lack- 
ing in  the  first  octave,  and  a  trivalent  [Ga]  and  a  quadrivalent  one  [Ge] 
in  the  second.  These  places  have  since  been  filled,  as  we  shall  pres- 
ently see.  The  first  two,  (the  short)  series  have  been  split  in  the  table, 
as  lithium  and  sodium  closely  resemble  potassium,  while  the  remaining 
members  of  these  series  fall  more  naturally  over  the  corresponding  ele- 
ments of  the  second  octave  of  the  third  series. 

The  fourth  series  (long)  is  nearly  complete.  It  begins  with  an 
active  alkali  metal,  rubidium,  and  ends  with  iodine,  a  halogen. 
The  rule  of  valence  is  strictly  preserved  throughout  the  series,  and  in 
general  the  elements  fall  below  those  which  they  most  closely  resemble. 

The  fifth,  sixth,  and  seventh  (long)  series  are  incomplete,  but  the 
order  of  the  atomic  weights  and  the  valence  enable  us  satisfactorily  to 
place  those  elements  which  are  known.  The  chemical  relations  to  ele- 
ments of  the  fourth  series  justify  the  position  assigned  to  each.  Cae- 
sium, for  example,  is  the  most  active  of  the  alkali  metals ;  barium  has 
always  been  classed  with  strontium,  and  bismuth  with  antimony. 

In  two  cases  a  slight  displacement  of  the  order  according  to  atomic 
weights  is  necessary.  Cobalt  is  put  before  nickel  because  it  resembles 
iron  more  closely.  Tellurium  and  iodine  are  placed  in  that  order  to 
bring  them  into  the  sulphur  and  halogen  groups  respectively.  Their 
valence  and  other  chemical  relations  both  require  this.  The  general 
agreement,  however,  is  very  remarkable. 

General  Relations  in  the  System.  —  In  every  octave  the  valence 
towards  oxygen  ascends  from  one  to  seven,  while  that  towards  hydro- 
gen, in  the  cases  of  the  four  last  elements  (when  they  combine  with 
hydrogen  at  all),  descends  from  four  to  one.  The  atomic  volume  disre- 
gards the  subsidiary  octaves  in  the  long  series.  It  descends  towards 
the  middle  of  each  series  (long  or  short),  and  ascends  again  towards 
its  initial  value.  The  other  physical  properties  fluctuate  within  the 
limits  of  each  series  in  a  similar  way.  The  values  of  each  physical 


410  INORGANIC   CHEMISTRY 

constant  for  corresponding  members  of  the  successive  series  do  not 
exactly  coincide,  however.  A  progressive  change,  as  we  descend  each 
vertical  column,  is  the  rule.  Thus  the  specific  gravities  (water  =  1) 
of  the  alkali  metals  rise  from  lithium  (0.59)  to  caesium  (1.85).  In  the 
same  group  the  melting-points  descend  from  lithium  (186°)  to  caesium 
(26.5°). 

It  must  be  stated  that  as  yet  no  exact  mathematical  relation  be- 
tween the  values  for  any  property  and  the  values  of  the  atomic  weights 
has  been  discovered  ;  only  a  general  relationship  can  be  traced.  An- 
ticipating the  discovery  of  some  more  exact  mode  of  stating  the  rela- 
tionship in  each  case,  and  remembering  that  similar  values  of  each 
property  recur  periodically,  usually  at  intervals  corresponding  to  the 
length  of  an  octave  or  series,  the  principle  which  is  assumed  to  un- 
derlie the  whole,  the  periodic  law,  is  stated  thus  :  All  the  properties  of 
the  elements  are  periodic  functions  of  their  atomic  weights. 

That  the  chemical  relations  of  the  elements  vary  just  as  do  the  phy- 
sical properties  of  the  simple  substances  is  easily  shown.  Thus,  each 
series  begins  with  an  active  metallic  (positive)  element,  and  ends  with 
an  active  non-metallic  (negative)  element,  the  intervening  elements 
showing  a  more  or  less  continuous  variation  between  these  limits. 
Again,  the  elements  at  the  top  are  the  least  metallic  of  their  respec- 
tive columns.  As  we  descend,  the  members  of  each  group  are  more 
markedly  metallic  (in  the  first  columns),  or,  what  is  the  same  thing, 
less  markedly  non-metallic  (in  the  later  columns ;  cf.  p.  403). 

In  the  first  series  boron  is  the  first  non-metal  we  encounter.  In  the 
second  series  silicon  is  the  first  such  element.  In  the  third  there  is 
more  difficulty  in  deciding.  Titanium,  vanadium,  and  germanium  are 
usually,  though  with  questionable  propriety,  classed  as  metals.  Sele- 
nium is  undoubtedly  a  non-metal.  Arsenic  is,  on  the  whole,  a  non- 
metal.  In  the  fourth  series  tellurium  is  commonly  considered  to  be 
the  first  non-metal.  Thus  a  zigzag  line,  shown  in  the  table,  separates 
all  the  non-metals  from  the  rest  of  the  elements,  and  confines  them  in 
the  right-hand  upper  corner. 

A  more  compact  form  of  the  table,  also  based  upon  one  of  Mendele- 
jefFs,  may  now  be  given  (Table  II).  The  only  difference  between 
this  and  the  other  is  that  the  two  octaves  of  each  long  series  have  been 
placed  in  the  same  set  of  seven  main  columns.  The  irregular  sets  of 
three  elements,  consisting  of  the  iron,  palladium,  and  platinum  groups, 
occupy  a  column  on  the  right  of  the  main  columns,  and  are  often 
called  collectively  the  eighth  group.  For  completeness,  the  newly 


SELENIUM  AND  TELLURIUM:  THE   PERIODIC   SYSTEM      411 


o 


3  S> 

*w   s* 


O      CO 


8    2 

II    « 

a 


a 


M 
II    S 


n  « 
S 


412  INORGANIC   CHEMISTRY 

discovered  elements,  found  chiefly  in  the  air,  have  been  placed  at  the 
left-hand  side.  Since  they  do  not  enter  into  combination  at  all,  their 
valence  may  appropriately  be  given  as  zero.  With  the  exception  of 
argon,  the  values  of  their  atomic  weights  agree  well  with  this  assign- 
ment. Hydrogen  is  the  only  element  whose  place  is  still  in  debate. 
Many  rare  elements  have  also  been  omitted,  and  tantalum  has  been 
placed  immediately  after  cerium.  The  valence  is  shown  by  the 
general  formulae  at  the  head  of  each  column. 

Applications  of  the  Periodic  System.  —  The  system  has  found 
application  chiefly  in  four  ways  : 

1.  By  leading  to  the  prediction  of  new  elements. 

2.  By  furnishing  a  comprehensive  classification  of  the  elements, 
arranging  them  so  as  to  exhibit  the  relationships  among  the  physical 
and  chemical  properties  of  the  elements  themselves  and  of  their  com- 
pounds.    Constant  use  will  be  made  of  this  property  of  the  table  in 
the  succeeding  chapters. 

3.  By  enabling  us  to  decide  on  the  correct  values  for  the  atomic 
weights  of  some  elements,  when  the   equivalent  weights  have   been 
measured,  but  no  volatile  compound  is  known  (cf.  pp.  201  and  212). 

4.  By  suggesting  problems  for  investigation. 

The  Prediction  of  New  Elements.  —  Mendelejeff  (1871)  drew 
attention  to  the  blank  then  existing  between  calcium  (40)  and  titanium 
(48).  He  predicted  that  an  element  to  fit  this  place  would  have  an 
atomic  weight  44  and  would  be  trivalent.  From  the  nature  of  the 
surrounding  elements,  he  very  cleverly  deduced  many  of  the  physical 
and  chemical  properties  of  the  unknown  element  and  of  its  com- 
pounds. He  named  it  eka-boron  (Skr.  eka,  one).  In  1879  Nilson 
discovered  scandium  (44),  and  its  behavior  corresponded  closely  with 
that  predicted  for  eka-boron. 

In  the  same  paper  Mendelejeff  described  two  other  elements, 
likewise  unknown  at  the  time.  They  were  to  occupy  vacant  places 
between  zinc  and  arsenic,  and  were  named  eka-aluminium  and  eka- 
silicon.  In  1875  Lecoque  de  Boisbaudran  found  gallium,  and  in  1888 
Winkler  discovered  germanium,  and  these  blanks  were  filled.  Other 
possible  elements,  such  as  eka-manganese  and  eka-antimony,  were 
described,  but  still  remain  to  be  discovered. 

Application  to  Fixing  Atomic  Weights.  —  Atomic  weights  have 
been  fixed  with  the  assistance  of  the  periodic  system  on  several  occasions. 


SELENIUM   AND   TELLURIUM  :   THE   PERIODIC   SYSTEM      413 

Thus  the  atomic  weight  of  uranium  was  thought  to  be  120  until  it  was 
observed  that  no  place  near  to  antimony  (120)  remained  unoccupied. 
With  the  value  240  (now  238.5),  the  element  was  accommodated  at 
the  foot  of  the  column  containing  those  which  it  most  resembled. 
Again,  the  equivalent  weight  of  indium  was  38,  and,  as  the  element  was 
supposed  to  be  bivalent,  it  received  the  atomic  weight  76.  It  was  quite 
but  of  place  near  arsenic  (75),  however,  being  decidedly  a  metal.  As  a 
trivalent  element  with  the  atomic  weight  115,  it  fell  between  cadmium 
and  tin.  Later  work  fully  justified  the  change.  Still  again,  glucinum, 
with  equivalent  weight  4.5,  resembled  aluminium  so  strongly  that 
it  was  thought  to  be  trivalent,  like  that  element,  and  to  have  the 
atomic  weight  13.5.  But  the  only  vacancy  in  the  first  series  then 
existing  was  between  lithium  (7)  and  boron  (11),  and  subsequent  in- 
vestigations showed  that  the  properties  of  glucinum  placed  it  most 
fittingly  in  that  position  as  a  bivalent  metal  with  the  atomic  weight  9. 
Finally,  and  quite  recently,  radium  (q.v.)  has  been  discovered,  and 
found  to  have  the  equivalent  weight  112.5  and  to  resemble  barium. 
If,  like  barium,  it  is  bivalent,  it  occupies  a  place  under  this  element, 
in  the  last  series. 

The  System  Suggests  Lines  of  Investigation.  —  The  periodic 
system  has  been  of  constant  service  in  the  course  of  inorganic  research, 
and  has  often  furnished  the  original  stimulus  to  such  work  as  well. 
For  example,  the  atomic  weights  of  the  platinum  metals  at  first  placed 
them  in  the  order,  Ir  (197),  Pt  (198),  Os  (199),  although  the  resem- 
blance to  iron  and  ruthenium  would  have  led  us  to  expect  that  osmium 
should  come  first.  For  similar  reasons  platinum  should  have  come  last, 
under  palladium.  A  reinvestigatioii  of  the  atomic  weights,  suggested 
by  these  considerations,  was  undertaken  by  Seubert,  and  the  old  values 
were  found  in  fact  to  be  very  inaccurate.  He  obtained :  Os  =  191, 
Ir  =  193,  Pt  =  195. 

Again,  the  atomic  weight  of  tellurium  bore  the  value  128  when  the 
table  was  first  constructed,  and  it  was  confidently  expected  that  re- 
examination  would  bring  this  value  below  that  of  iodine  (then  127, 
now  126.97).  The  problem  has  proved  more  difficult  than  usual,  and 
several  most  careful  studies  of  the  subject  have  been  made  by  chemists, 
using  different  methods.  Absolute  certainty  has  not  yet  been  reached, 
but  it  seems  probable  that  the  real  value  of  the  atomic  weight  is  not 
far  from  Te  =  127.6,  and  therefore  more  than  half  a  unit  greater  than 
that  of  iodine.  Since,  however,  mathematical  correspondence  is  found 


414  INORGANIC   CHEMISTRY 

nowhere  in  the  system,  the  existence  of  marked  inconsistencies  like 
this  need  not  shake  our  confidence  in  its  value  when  it  is  used  with 
due  consideration  of  the  degree  of  correspondence  to  be  expected. 

Originally  lead,  although  it  fell  in  the  fourth  column,  possessed 
only  one  compound,  Pb02,  in  which  it  seemed  to  be  undoubtedly  quad- 
rivalent. Search  for  salts  of  the  same  form,  however,  speedily  yielded 
the  tetrachloride  (PbCl4),  tetracetate,  and  many  others.  As  has  been 
stated  already,  there  are  many  elements  between  lanthanum  and 
tantalum,  about  which  sufficient  knowledge  has  not  yet  been  accumu- 
lated to  make  possible  their  definite  allotment  to  places  in  the  table. 
The  existence  of  osrnic  acid  (Os04),  and  a  corresponding  compound 
of  ruthenium,  suggests  that  other  compounds  of  the  elements  of  the 
eighth  group,  displaying  the  valence  eight,  may  be  capable  of  prepa- 
ration. The  collocation  of  copper,  silver,  and  gold,  in  the  same  column 
with  the  alkali  metals,  is  not  at  present  perfectly  satisfactory,  and  sug- 
gests the  advisability  of  strengthening  their  position,  if  possible,  by 
further  investigation. 

The  System  as  a  Guide  in  Classification.  — Having  disposed 
of  the  halogen  and  sulphur  families,  situated,  respectively,  in  the 
seventh  and  sixth  columns  of  the  table,  we  shall  next  take  up  nitrogen 
and  phosphorus  from  the  right  side  of  the  fifth  column.  This  family 
includes  nitrogen,  phosphorus,  arsenic,  antimony,  and  bismuth,  and 
brings  us  for  the  first  time  in  contact  with  a  group  which  is  partly 
non-metallic  and  partly  metallic.  Then  from  the  fourth  column,  we 
shall  select  carbon  and  silicon,  and  from  the  third  boron,  leaving  the 
other,  more  decidedly  metallic  elements  for  later  treatment. 

Exercises.  —  1.  Can  you  explain  the  presence  of  free  selenium  in 
the  flues  of  pyrite-burners  (p.  385)  ? 

2.  Why  does  the  existence  of  tellurium  tetrachloride  in  solution 
in  aqueous  hydrochloric  acid  show  tellurium  to  be  somewhat  metallic 
in  chemical  properties  ? 

3.  How  should  you  attempt  to  obtain  HIO  and  HBr04,  or  their 
salts  ? 

4.  Make  a  list  of  bivalent  elements  and  criticize  this  method  of 
grouping  as  a  means  of  chemical  classification. 

5.  Write  down  the  symbols  of  the  elements  in  the  fourth  series 
of  Table  I  (that  beginning  with  rubidium,  and  ending  with  iodine). 
Record  the  valence  of  each  element  toward  oxygen,  using  for  reference 
the  chapters  in  which  the  oxygen  compounds  are  described. 


CHAPTER  XXIV 
NITROGEN  AND   ITS    COMPOUNDS    WITH    HYDROGEN 

WHEN  the  oxygen  of  the  air  is  removed  (p.  62),  a  gas  remains 
which  is  largely  nitrogen.  When  first  discovered,  the  most  conspicu- 
ous property  this  gas  was  observed  to  have  was  indifference  ;  it  did  not 
support  combustion  or  life.  The  latter  fact  led  to  its  being  named 
azote  (Grk.  £<OTIKOS,  life).  While  this  name  is  still  preserved  in  the 
French  language,  in  English  the  name  of  the  substance  is  derived 
from  the  fact  that  it  is  an  important  constituent  of  saltpeter  KN08 
(Lat.  nitrum). 

The  Chemical  Relations  of  the  Element.  —  In  some  compounds 
nitrogen  is  trivalent,  while  in  others,  particularly  those  containing 
oxygen  and  other  negative  elements,  it  is  quinquivalent. 

The  compounds  of  nitrogen  are  often  extremely  active  and  inter- 
esting. Those  of  them  which  we  have  to  discuss  in  inorganic  chem- 
istry are  ammonia  (NH3)  and  nitric  acid  HN03,  and  several 
related  substances.  The  organic  compounds  containing  nitrogen  are 
very  numerous  and  possess  highly  characteristic  properties.  Some, 
like  nitroglycerine  ( q.  v.)  and  gun-cotton,  are  violently  explosive ; 
others,  like  antipyrine,  show  great  physiological  activity ;  still  others, 
such  as  the  aniline  and  other  organic  dyes,  provide  us  with  beautiful 
and  useful  coloring  matters. 

Occurrence.  —  Apart  from  the  presence  of  free  nitrogen  in  the 
air,  the  element  is  found  in  many  forms  of  combination.  The  nitrates 
of  potassium  and  sodium  are  found  in  Bengal  and  Peru  respectively. 
Natural  manures,  such  as  guano,  contain  large  quantities  of  nitrogen 
compounds,  and  owe  part  of  their  value  as  fertilizers  to  this  fact. 
Nitrogen  is  an  essential  constituent  of  vegetable  and  animal  matter. 
The  albumins,  for  example,  which  often  make  up  a  considerable  part 
of  such  matter,  contain  on  an  average  about  15  per  cent  of  combined 
nitrogen. 

416 


416  INORGANIC   CHEMISTRY 

Preparation.  —  Nitrogen  containing  about  one  per  cent  of  argon 
(q.v.)  is  easily  obtainable  from  purified  air  when  the  oxygen  of  the 
latter  is  removed.  For  this  purpose  phosphorus  is  frequently  burned 
in  the  air,  or  the  air  is  passed  over  heated  copper. 

When  pure  nitrogen  is  required,  it  must  be  obtained  from  chemical 
compounds,  in  order  that  it  may  be  free  from  argon.  The  simplest 
method  is  to  heat  ammonium  nitrite.  This  substance  decomposes 
readily,  even  at  temperatures  little  above  the  ordinary,  according  to 
the  following  equation  : 


In  practice,  since  ammonium  nitrite  cannot  easily  be  kept,  a  mixture 
of  an  ammonium  salt  with  some  salt  of  nitrous  acid  is  employed. 
Thus,  when  strong  solutions  of  ammonium  chloride  and  sodium  nitrite 
are  mixed,  a  double  decomposition  results  in  the  formation  of  ammo- 
nium nitrite,  NH4C1  +  NaN02  <±  NH4N02  +  NaCl,  and  this  breaks 
up  when  heat  is  applied,  giving  nitrogen. 

In  certain  cases  it  is  convenient  to  prepare  nitrogen  by  the  oxida- 
tion of  ammonia  (NH3),  by  passing  the  latter  over  heated  cupric  oxide, 
or  by  the  reduction  of  nitric  oxide  (NO)  by  passing  this  gas  over  heated 
copper 

Physical  Properties.  —  Nitrogen  is  a  colorless,  tasteless,  odorless 
gas,  as  we  should  expect  from  the  fact  that  air  possesses  these  proper- 
ties. When  sufficiently  cooled,  it  forms  a  colorless  liquid,  boiling  at 
—  194°.  By  further  cooling,  this  liquid  freezes  to  a  white  solid,  which 
melts  at  —214°.  One  liter  of  pure  nitrogen  weighs  1.2507  grams. 
The  solubility  of  nitrogen  in  water  :  1.6  volumes  in  100,  is  less  than 
that  of  oxygen. 

Chemical  Properties.  —  The  density  of  the  gas  shows  the  formula 
of  free  nitrogen  to  be  N2. 

Nitrogen  unites  with  few  common  chemical  elements  directly.  At 
ordinary  temperatures  it  is  almost  absolutely  indifferent.  When 
passed  through  a  tube  over  strongly  heated  lithium,  calcium,  magne- 
sium, or  boron,  it  forms  definite  chemical  compounds,  known  as 
nitrides,  in  which  it  is  trivalent.  These  have  the  formulae  Li3N,  Ca3N2, 
Mg3N2,  and  BN  respectively.  When  the  gas  is  mixed  with  oxygen  or 
hydrogen,  and  the  mixture  is  heated,  little  chemical  action  takes  place. 
When  sparks  from  an  induction  coil  are  passed  between  platinum 


NITROGEN   AND   ITS   COMPOUNDS    WITH   HYDROGEN         417 

wires  through  the  mixtures,  small  amounts  of  nitrogen  tetroxide 
N204,  and  ammonia  NH8,  respectively,  are  produced.  The  indifference 
of  free  nitrogen  is  doubtless  due  to  the  fact  that  its  molecules  (N2)  are 
extremely  stable. 

One  case  of  direct  union  of  nitrogen  is  of  economic  importance. 
The  supply  required  by  plants  is  obtained  partly  from  nitrogen  com- 
pounds contained  in  fertilizers,  or  equivalent  substances  already  pres- 
ent in  the  soil,  and  partly  from  ammonium  nitrite  and  nitrate,  which 
are  washed  down  from  the  air  by  the  rain.  It  appears,  however,  that 
plants  belonging  to  the  order  leguminosce,  such  as  peas,  beans,  clover, 
etc.,  may  be  associated  constantly  with  certain  bacteria  which  nourish 
in  nodules  upon  their  roots.  These  bacteria  have  the  power  of  taking 
free  nitrogen  from  the  air,  which  penetrates  the  soil,  breaking  up  its 
molecules,  and  producing  compounds  containing  nitrogen.  The  masses 
round  the  roots  often  contain  over  five  per  cent  of  combined  nitrogen 
which  has  been  acquired  in  this  way.  These  compounds  are  chiefly 
albumins,  which  are  afterwards  digested  and  absorbed  by  the  roots  of 
the  plant. 

Compounds  of  Nitrogen  and  Hydrogen.  —  The  commonest  and 
longest  known  of  these  substances  is  ammonia  NH3,  which  was  first 
described  by  Priestley  (1774)  and  named  "  alkaline  air."  Curtius 
discovered  hydrazine  N2H4  in  1889,  and  hydrazoic  acid  HN8  in  1890. 
Hydroxylamine  NH30,  discovered  by  Lossen  in  1865,  is  similar  to 
ammonia  in  chemical  behavior. 

AMMONIA. 

Ammonia  is  liberated  in  connection  with  the  putrefaction  of  animal 
matter. 

Preparation.  —  1.  When  sparks  from  an  induction  coil  are  passed 
through  a  mixture  of  nitrogen  and  hydrogen,  small  proportions  of  the 
materials  unite  to  form  ammonia  (see  below). 

2.  When  water  is  added  to  the  nitride  of  magnesium  or  calcium 
(p.  416),  ammonia  is  given  off  and  the  hydroxide  of  the  metal  formed  : 

Mg8N2  +  6H20  ->  3Mg(OH)2  +  2NH3. 

3.  When  parts  of  animals,  particularly  the  horns,  hides,  and  feathers, 
which  contain  complex  compounds  of  carbon,  nitrogen,  hydrogen,  and 
oxygen,  are  heated  strongly,  much  of  the  nitrogen  is  driven  out  as 


418  INORGANIC   CHEMISTRY 

ammonia.  Hence  the  name  spirit  of  hartshorn,  applied  to  the  aqueous 
solution. 

4.  Coal  often  contains  as  much  as  2  per  cent  of  combined  nitrogen, 
and  when  heated  in  absence  of  air,  gives  off  much  ammonia.     The 
entire  commercial  supply  is  obtained  as  a  by-product  from  operations 
like  the  manufacture  of  illuminating-gas  and  of  coke,  in  which  the 
destructive  distillation  of  coal  takes  place.     The  crude  mixture  of  gases 
passes  first  through  water,  in  which  some  of  the  tar  is  condensed  and 
most  of  the  ammonia  dissolves.     This  ammoniacal  liquor  is  heated  with 
a  little  slaked  lime,  and  the  escaping  gas  is  led  into  dilute  hydrochloric 
or  sulphuric  acid,  with  which  it  combines  to  form  the  chloride  or  sul- 
phate of  ammonium. 

5.  In  the  laboratory,  a  mixture  of  slaked  lime  and  some  salt  of 
ammonium,  such  as  ammonium  chloride,  either  with  or  without  water, 
is  heated  in  a  flask  or  retort  provided  with  a  delivery  tube  : 

Ca(OH)2  +  2NH4C1  <=»  CaCL,  +  2NH4OH, 
NH4OH<r±H20+NH8. 

By  the  double  decomposition  usual  with  ioiiogens,  ammonium  hydroxide 
is  formed,  and  this,  being  unstable,  immediately  breaks  up  into  water 
and  ammonia. 

6.  Warming  the  aqueous  solution  also  gives  a  steady  stream  of  the 
gas.     Since  the  gas  is  very  soluble  in  water,  it  must  be  collected  over 
mercury  or  by  upward  displacement  of  air.     The  moisture  which  it  may 
contain  is  removed  by  passage  through  a  tower  or  wide  tube  filled  with 
quicklime.     Calcium  chloride  cannot  be  used,  as  the  gas  combines  with 
it  forming  a  compound  CaCl2,  8NH3,  similar  in  properties  to  a  hydrate 
(pp.  120-123). 

Physical  Properties.  —  Ammonia  is  a  colorless  gas  with  a  pun- 
gent, characteristic  odor  familiar  in  smelling-salts.  The  G.M.V.  of  the 
gas  weighs  17.29  g.,  so  that  the  density  is  little  more  than  half  that  of 
air  (cf.  p.  215)  When  liquefied  it  boils  at  —  34°  and  exercises  a  pressure 
of  6.5  atmospheres  at  10°.  The  solid  is  white  and  crystalline  and  melts  at 
—  77°.  The  gas  is  very  soluble  in  water,  one  volume  of  the  latter  dissolv- 
ing 1148  volumes  of  the  gas  at  0°,  764  volumes  at  16°,  and  306  volumes 
at  50°.  The  35  per  cent  aqueous  solution,  saturated  at  15°,  and  sold 
as  "  concentrated  ammonia,"  has  a  sp.  gr.  0.882.  The  whole  of  the 
dissolved  gas  may  be  removed  by  boiling  (cf.  p.  379). 


NITROGEN  AND   ITS   COMPOUNDS   WITH   HYDROGEN         419 


Liquefied  ammonia  is  used  in  refrigeration.  The  gas  is  liquefied  by 
compression,  and  the  heat  which  is  thus  liberated  is  removed  by  flowing 
water  which  surrounds  the  pipes.  The  liquid  then  flows  into  other 
pipes  immersed  in  calcium  chloride  brine,  and  is  there  allowed  to  evapo- 
rate, the  gas  returning  to  the  compressor.  The  heat  of  vaporization, 
260  calories  per  gram,  is  taken  from  the  brine,  which  is  thus  partially 
frozen.  The  resulting  freezing  mixture  of  ice  and  calcium  chloride 
solution  is  then  distributed  to  the  localities  to  be  cooled.  The  ammonia 
and  brine  remain  within  their  respective  closed  systems  of  pipes, 
and  are  used  over  and  over  again.  In  many  cases  the  cooling  by  the 
liquefied  ammonia  is  used  directly,  and  the  intermediate  freezing  mix- 
ture is  dispensed  with. 

Chemical  Properties.  —  The  discharge  from  an  induction-coil 
decomposes  ammonia  (to  the  extent  of  94-98  per  cent)  into  nitrogen 
and  hydrogen.  Under  the  same  circumstances,  union  of  the  constitu- 
ents also  occurs  (up  to  2-6  per  cent)  : 


The  behavior  of  a  system  in  chemical  equilibrium  may  be  illustrated  by 

inclosing  dry  ammonia  in  a  tube  over  mercury  and  allowing  sparks  to 

pass  between  platinum  wires  (Fig.  85).     After  a  time 

the  volume  is  practically  doubled,  every  two  molecules 

of   ammonia   giving  four  of  the  products  (cf.  p.  208). 

When,  now,  a  little   sulphuric  acid  is  admitted  above 

the  mercury,  the  remaining  trace  of  ammonia  combines 

with  the  acid  and  is  removed.     This  disturbs  the  equil- 

ibrium.     As  sparks   continue  to  pass,  the  action  now 

proceeds  steadily  backwards.      The  ammonia  which  is 

formed  is  no  longer  subject  to  decomposition,  for,  once 

it  has  combined  with  the  acid,  it  never  returns  to  the 

neighborhood  of  the  discharge.     Thus  the  whole  of  the 

gases  finally  combine  : 


FIG.  85. 


-  H2S04  +  2KH8  *=>  N2  +  3H,, 

and  ammonium  sulphate,  dissolved  in  the  excess  of  acid, 
alone  remains.     The  action,  therefore,  first  goes  almost 
completely  in  one  direction,  and  then  quite  completely  in  the  other, 
while  no  change  has  taken  place  in  the  conditions  to  which  the  gas 
is  subjected  at  the  point  where  the  interaction  is  occurring.     The  sole 


420  INORGANIC   CHEMISTRY 

difference  is  that  a  little  of  an  acid  has  been  introduced  into  a  rela- 
tively remote  part  of  the  space. 

When  passed  over  reducible  oxides,  such  as  heated  cupric  oxide, 
ammonia  is  oxidized  to  water  and  free  nitrogen  : 

3CuO  +  2NH3  ->  3Cu  +  3H2O  +  N2, 

and  burns  in  pure^  oxygen  with  the  same  result.  In  air,  heat  is  used 
up,  not  only  in  decomposing  the  compound,  but  also  in  raising  the 
temperature  of  the  nitrogen  of  the  atmosphere.  This  causes  a  contin- 
uous drain  on  the  heat  given  out  by  the  action,  and  prevents  the  main- 
tenance of  the  kindling  temperature  (p.  73). 

When  dry  ammonia  is  passed  over  heated  potassium  or  sodium,  one 
unit  of  hydrogen  is  displaced : 

2K  +  2NH3  ->  2KNH2  +  H2, 

and  a  solid  of  metallic  appearance  remains  behind.  Substances  con- 
taining the  group  —  NH2  are  often  called  amides,  and  this  is  therefore 
named  potassamide. 

Chlorine  and  bromine  combine  with  the  hydrogen  and  liberate  the 
nitrogen  of  ammonia.  This  action  may  be  used  for  obtaining  a  stream 
of  nitrogen,  provided  excess  of  chlorine  is  avoided  (see  Nitrogen  tri- 
chloride, below).  Chlorine  is  led  into  a  solution  of  ammonium  chloride : 

2NH4C1  +  3C12  ->  N2 1  +  8HC1. 

The  most  characteristic  property  of  ammonia  is  its  power  to  form 
a  base  by  combination  with  water : 

NH8  (gas)  fc?  NH8  (diss'd)  +  H2  0  <=±  NH4OH  ±=»  NH4*  +  OH'. 

Probably  only  a  small  proportion  of  the  gas  is  actually  combined  at 
any  one  time,  the  greater  part  being  simply  dissolved. 

The  gas  unites  also  with  acids,  forming  salts  (cf.  p.  337),  which,  in 
solution,  are  highly  ionized  :, 

NH3    +    HC1-»NH4C1, 
NH8  +  HNO3  ->NH4N03. 

Ammonium  Compounds.  —  Since  NH4  plays  the  part  of  a  metal, 
entering  into  the  composition  of  a  base  and  of  a  series  of  salts,  it  is 
named  ammonium.  It  constitutes  the  positive  ion  of  these  compounds. 
As  this  radical  forms  a  univalent  ion  and  gives  a  distinctly  alkaline 
base,  it  is  classed  with  potassium  and  sodium  as  one  of  the  metals  of 
the  alkalies  .v.. 


NITROGEN  AND   ITS   COMPOUNDS   WITH    HYDROGEN         421 

Ammonium  hydroxide,  although  much  less  completely  ionized 
than  potassium  hydroxide,  affects  litmus  easily.  In  a  normal  solution 
about  0.4  per  cent  of  the  ammonia  is  in  the  form  of  ammonion, 
NH4*.  When  an  acid  is  added  to  the  solution,  the  correspondingly 
small  amount  of  hydroxidion  which  exists  in  it  is  removed  and  the 
various  equilibria  are  displaced  forwards.  The  final  result  is  the  same 
as  with  any  other  base  : 


NH8  (diss'd)  +  H20  <r±  NH4OH  <          4 

ci'  +H- 


When  strongly  heated,  all  ammonium  salts  are  decomposed  and,  usually, 
give  ammonia  and  the  acid.  When  the  latter  is  volatile,  the  whole 
material  of  the  salt  is  thus  converted  into  gas.  If  the  acid  is  volatile  with- 
out permanent  decomposition,  it  reunites  with  the  ammonia  when  the 
vapor  is  cooled  : 

NH4C1  <r±  HCI  +  NH8. 

This  behavior  distinguishes  ammonium  salts  from  those  of  the  typical 
metals,  for,  with  the  exception  of  mercury  salts,  most  other  salts  are  not 
easily  and  completely  volatilized.  The  use  of  ammonium  chloride  (sal- 
ammoniac)  in  soldering  depends  on  the  dissociation  of  the  salt,  by  the 
heat  of  the  iron,  and  the  action  of  the  liberated  hydrochloric  acid  on 
the  oxide  which  covers  the  surface  of  the  metal  to  be  soldered. 

Ammonium  salts  are  recognized  by  warming  them,  dry  or  in  solu- 
tion, with  a  base  : 


when  the  odor  of  ammonia  becomes  noticeable.  When  the  solution  is 
used,  it  is  the  tendency  of  the  NH4*  and  OH'  to  unite  to  form  the  slightly 
ionized  molecular  hydroxide  that  sets  the  other  equilibria  in  motion. 
The  principle  at  the  basis  of  the  change  is  thus  the  same  as  in  neu- 
tralization (p.  354). 

In  ammonia,  nitrogen  is  trivalent,  while  in  the  salts  it  appears  to 
be  quinquivalent  : 

H\  H\         H         H\         H 

H-N       H-NX  H-NX 

H/  H/      NOH       H/     ^C\ 


422  INORGANIC   CHEMISTRY 


HYDRAZINE,  HYDRAZOIC  ACID,  HYDROXYLAMINE. 

Hydrazine.  —  By  reduction  of  a  compound  of  nitric  oxide  and 
potassium  sulphite  by  means  of  sodium  amalgam,*  a  solution  of  hydra- 
zine hydrate  is  obtained : 

K^SOg,  2KO  +3H2  _>  N2H4,  H2O  +  K2S04. 

The  same  substance  is  more  easily  made  from  certain  organic  deriva- 
tives. When  the  hydrate  is  distilled  with  barium  oxide,  under  reduced 
pressure,  hydrazine  is  liberated : 

N2H4,  H20  +  BaO  ->  N2H4|  +  Ba(OH)2. 

Hydrazine  is  a  white  solid,  which  fumes  in  moist  air,  giving  the 
hydrate  once  more.  It  melts  at  1.4°  and  boils  at  113.5°. 

Hydrazine  hydrate  freezes  at  about  —40°,  boils  at  118.5°,  and  can 
be  distilled  without  decomposition.  Its  aqueous  solution  is  alkaline, 
and  salts  can  be  formed  by  neutralization. 

Hydrazoic  Acid.  —  When  nitrous  oxide  (q.v.)  is  led  over  sodamide 
at  200°,  water  is  liberated  and  sodium  hydrazoate  remains  behind : 


H2O. 

A  dilute  solution  of  the  free  acid  is  best  obtained  by  distilling  the  lead 
salt  with  dilute  sulphuric  acid.  A  similar  solution  may  be  made  more 
directly  by  adding  cold  nitrous  acid  (q.v.)  to  a  cold  aqueous  solution 
of  hydrazine  hydrate  : 

N2H6OH  +  HN02  -»  HN8  +  3H20. 


By  repeated  distillation  of  the  solution  the  pure  acid  is  obtainable.  It 
boils  at  37°.  The  operation  is  a  dangerous  one,  as  the  pure  acid  is 
violently  explosive,  resolving  itself  into  nitrogen  and  hydrogen  with 
liberation  of  much  heat  : 

HN8,  Aq  ->  H  +  3K  +  Aq  +  61,600  cal. 

It  is  an  acid  of  somewhat  greater  activity  and  degree  of  ionization  than 
acetic  acid.  Active  metals,  like  magnesium,  displace  hydrogen  from 
its  solution.  Its  silver  salt  (AgN3)  is  insoluble  in  water.  It  neutral- 

*  The  sodium  dissolved  in  the  mercury  interacts  with  the  water,  giving  hydro- 
gen (see  Active  state  of  hydrogen,  below). 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN        423 

izes  ammonium  hydroxide  and  hydrazine  hydrate,  giving  two  salts, 
NH4N8  and  N2H5N3.  These  constitute  two  additional  compounds  of 
nitrogen  and  hydrogen,  but  differ  from  ammonia  and  hydrazine  in 
being  ionogens. 

Hydroxylamine.  —  Tin  displaces  hydrogen  from  dilute  hydro- 
chloric acid  :  Sn  -j-  2HC1  — >  SnCLj  +  H2,  and  this  combination  forms 
a  reducing  agent  (see  below).  When  dilute  nitric  acid  (q.v.)  is  added 
to  the  mixture,  a  considerable  part  of  it  is  reduced  to  hydroxylamine  : 

HN03  +  3H2->  NH80  +  2H20. 

The  hydroxylamine  forms  a  weak  base,  NH4O.OH,  with  water  which 
interacts  with  the  excess  of  acid,  giving  hydroxylamine  hydrochloride, 
NH4OC1.  By  more  complete  reduction  of  part  of  the  nitric  acid,  some 
ammonium  chloride  is  formed  at  the  same  time.  To  secure  the  salt 
of  hydroxylamine,  the  tin  ions  are  removed  by  means  of  hydrogen  sul- 
phide, which  precipitates  stannous  sulphide.  The  filtered  solution  is 
then  evaporated  to  dryness,  the  hydroxylamine  hydrochloride  is 
extracted  from  the  residue  with  absolute  alcohol,  and  this  alcoholic 
solution  is  finally  evaporated  in  turn.  The  hydrochloride  is  a  white 
crystalline  salt. 

When  the  hydrochloride  is  treated  with  a  base,  in  the  absence  of 
water,  and  the  mixture  is  distilled  under  reduced  pressure,  hydroxyla- 
mine passes  over.  It  is  a  white  solid  melting  at  33°  and  boiling  at 
58°  at  22  mm.  pressure.  Even  before  melting  (above  15°),  it  begins  to 
decompose,  and  explodes  at  or  below  130°.  In  chemical  behavior  it  is 
like  ammonia.  With  water  it  forms  a  base  which  combines  with  acids, 
but  is  less  active  than  ammonium  hydroxide.  It  is  a  stronger  reducing 
agent  than  ammonia,  precipitating  silver  from  a  solution  of  silver  nitrate. 

The  union  with  acids  indicates  that  the  molecule  of  hydroxylamine 
is  unsaturated,  and  hence  the  nitrogen  unit  is  supposed  to  be  trivalent : 

H\  H\         H 

H-N  H-NX 

H-0/  H-0/      XC1 

An  Active  State  of  Hydrogen  (Nascent  Hydrogen).  —  Pure 

hydrogen  gas,  whatever  its  source  may  have  been,  shows  conspicuous 
reducing  powers  only  when  heated  (cf.  p.  109).  But  when  the  gas  is 
absorbed  in  platinum  or  palladium  (p.  107),  the  two  together  reduce 
many  substances  in  the  cold,  although  the  platinum  and  palladium 
themselves  seem  to  take  no  part  in  the  change.  They  are  simply 


424  INORGANIC   CHEMISTRY 

catalytic  or  contact  agents,  and  in  their  presence  the  hydrogen  is  more 
active.  Similarly  the  hydrogen  liberated  by  the  electrolysis  of  a  dilute 
acid  reduces  substances  ad'ded  to  the  liquid  surrounding  the  cathode 
(cf.  p.  304),  and  shows  different  degrees  of  activity  according  to  the 
material,  platinum,  carbon,  or  the  like,  of  which  the  cathode  is  made. 
Finally,  zinc,  tin,  and  other  metals  with  dilute  acids  give  hydrogen, 
and  this  can  produce  reductions  so  long  as  it  is  in  the  immediate 
neighborhood  of  the  metal  (cf.  pp.  422,  423,  and  see  Arsine).  Thus,  as 
we  have  just  seen,  nitric  acid  is  reduced  to  hydroxylamine  when  it  is 
mixed  with  the  materials  which  generate  the  hydrogen.  The  hydrogen 
on  the  platinum  plate  forming  the  cathode  of  an  electrolytic  cell  is 
equally  able  to  reduce  nitric  acid.  The  very  same  hydrogen,  however, 
if  led  by  a  tube  from  a  Kipp's  apparatus  or  electrolytic  cell  into  a 
second  vessel  would  be  then  quite  inactive  and  incapable  of  affecting 
the  same  nitric  acid.  Hydrogen,  therefore,  like  other  substances,  may 
show  a  greatly  increased  speed  of  interaction  when  in  immediate  con- 
tact with  a  suitable  body.  This  more  active  state  of  hydrogen  is  de- 
scribed as  the  nascent  state,  because  it  happens  to  be  a  common  con- 
dition of  hydrogen  when  associated  with  substances  which  produce  it. 
This  state  has,  however,  no  necessary  connection  with  such  an  imme- 
diately preceding  act  of  liberation,  as  our  first  example  shows. 

Some  theorists  apply  the  atomic  hypothesis  to  this  phenomenon,  and  suggest 
that  it  is  the  freshly  liberated  atomic  hydrogen  (H)  which  possesses  this  greater 
activity.  This  would  require  us  to  believe  that  nascent  hydrogen  was  a  different 
substance  from  free  hydrogen,  which  may  or  may  not  be  true,  but  is  certainly  im- 
probable. If  the  hypothesis  of  atomic  hydrogen  were  correct,  all  arrangements  that 
give  free  hydrogen  under  similar  circumstances  (e.g.  electrolytic  cells  with  different 
electrode-materials)  should  show  equal  reducing  powers.  But  it  is  notorious  that 
they  do  not.  Different  substances  are  known,  however,  to  show  different  efficiencies 
as  contact  agents. 

The  atomic  theory  of  nascent  action  is  often  used  to  explain  the  oxidizing 
activity  of  hypochlorous  acid  (cf.  p.  271).  But  it  is  less  often  employed  to  explain  the 
oxidizing  power  of  sulphuric  acid  or  other  oxidizing  agents.  Yet,  logically,  it 
ought  to  be  applied  even  to  double  decomposition,  if  it  is  used  at  all.  If  this 
were  done,  salt  and  sulphuric  acid  would  be  assumed  to  give  nascent  chlorine  and 
nascent  hydrogen,  respectively,  since  they  yield  hydrogen  chloride  in  the  cold, 
while -free  hydrogen  and  chlorine  do  not. 

Halogen  Compounds  of  Nitrogen.  —  When  ammonium  chloride 
solution  is  treated  with  excess  of  chlorine,  drops  of  an  oily  liquid,  nitro- 
gen trichloride,  are  formed : 

3C12  +  NH4C1  ->  NC13  +  4HC1. 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN        425 

It  is  extremely  explosive,  resolving  itself  into  its  constituents  with 
liberation  of  much  heat. 

There  appear  to  be  several  compounds  related  to  ammonia  and  con- 
taining iodine.  When  a  solution  of  iodine  in  potassium  iodide  solu- 
tion (p.  156)  is  added  to  aqueous  ammonia,  a  brown  precipitate  is 
formed.  This  seems  to  have  the  composition  N2H8I3,  and  is  named  nitro- 
gen iodide.  It  may  be  handled  while  wet,  but  when  dry  decomposes  into 
its  constituents  with  violent  explosion  if  touched  with  a  feather. 

Exercises.  — 1.  When  moist  air  is  used  as  a  source  of  nitrogen,  what 
advantage  is  there  in  using  copper  rather  than  the  less  expensive 
metal  iron,  for  removing  the  oxygen  (p.  416)  ? 

2.  How  many  grams  of  water  at  0°  could  be  frozen  (p.  115)  by  the 
removal  of  the  heat  required  to  evaporate  50  g.  of  liquid  ammonia 
(p.  419)  ? 

3.  How  many  grams  of  ammonia  are  contained  in  1  1.  of  "  concen- 
trated ammonia"  (p.  418)  ? 

4.  What  are  the  ions  of  hydrazine  hydrate  (p.  422)  ?     Formulate 
(p.  354)  the  neutralization  of  this  base  with  sulphuric  acid. 

5.  What  is  the  object  attained  by  distilling  under  reduced  pressure 
in  making  hydrazine  (p.  422)  arid  hydroxylamine  (p.  423)  ? 

6.  Classify  (p.  187)  the  interaction  of  a  nitride  with  water  (p.  417), 
and  of  chlorine  and  ammonium  chloride  (p.  420),  and  the  results  of 
heating  ammonium  nitrite  (p.  416)  and  ammonium  chloride  (p.  421). 


CHAPTER   XXV 
THE   ATMOSPHERE.      THE   HELIUM  FAMILY 

WE  have  seen  that  to  counterbalance  the  pressure  of  the  air  a 
column  of  mercury  of  the  same  diameter  averaging  760  mm.  in  height 
is  required.  Let  th'e  section  of  the  column  be  1  sq.  cm.  Then  the 
pressure  of  a  column  of  air  1  sq.  cm.  in  section,  and  extending  so 
far  from  the  earth  as  any  downward  tendency  of  the  air  exists,  is 
equal  to  the  weight  of  a  column  of  mercury  containing  76  c.c.  of  the 
metal.  The  weight  of  1  c.c.  of  mercury  being  13.6  g.,  this  volume  of 
the  metal  weighs  1033.6  g.  This  number  represents  therefore  the 
pressure  which  is  exerted  by  the  air  upon  each  square  centimeter  of 
the  earth's  surface.  In  ordinary  units  of  measure,  this  is  nearly  fif- 
teen pounds  to  the  square  inch. 

A  more  vivid  appreciation  of  the  reality  of  this  pressure  may  be  obtained  by 
noticing  one  of  its  effects.  By  boiling  a  small  quantity  of  water  in  a  tin  can 
furnished  with  a  narrow  opening,  we  remove  the  whole  of  the  air  from  its  interior, 
displacing  it  by  steam.  While  the  boiling  is  in  progress,  we  suddenly  close  the 
opening  with  a  tightly  fitting  cork  and  remove  the  burner.  While  the  steam  was 
still  issuing  from  the  opening,  its  pressure  was  practically  that  of  the  atmosphere, 
and  the  can  was  subject  to  the  same  pressure  inside  and  out.  With  the  removal 
of  the  flame,  however,  the  steam  condenses,  and  the  pressure  on  the  interior  is 
reduced  to  a  minute  fraction  of  its  original  value,  while  the  pressure  on  the 
exterior  is  still  the  same  (1  atmosphere).  Under  this  pressure  a  vessel  of  ordinary 
tin-plate  completely  collapses. 

Components  of  the  Atmosphere.  —  The  first  component  of  the 
air  to  be  recognized  and  studied  was  oxygen.  When  a  rusting  metal  or 
burning  body  of  any  kind  removes  the  oxygen  by  combining  with  it,  a 
gas  remains  which  represents  about  four-fifths  of  the  original  volume. 
This  residual  gas  is  mainly  nitrogen.  It  contains,  however,  small  pro- 
portions of  several  inert  gases,  of  which  the  most  plentiful  is  argon. 
Examination  before  the  removal  of  the  oxygen  also  reveals  the  presence 
of  varying  proportions  of  carbon  dioxide,  water  vapor,  and  ammonium 
nitrate.  In  the  neighborhood  of  cities  the  air  likewise  contains  sulphur 
dioxide,  hydrogen  chloride,  hydrogen  sulphide,  and  other  ingredients 
which  may  be  described  as  accidental.  There  are  thus  three  classes 

426 


THE  ATMOSPHERE.     THE   HELIUM  FAMILY  427 

of  substances  in  the  air.  Those  of  the  first  class,  oxygen,  nitrogen,  and 
argon,  are  present  in  almost  constant  quantities ;  those  of  the  second 
class  are  very  variable  in  quantity,  although  found  in  all  samples  of 
air ;  those  of  the  third  class  are  accidental.  Finally  j  one  significant 
component  of  the  air  is  the  dust  which  a  powerful  beam  of  light  reveals 
as  it  passes  through  the  atmosphere  of  a  darkened  room. 

Components  which  are  Constant  in  Amount.  —  The  determi- 
nation of  the  oxygen  by  burning  phosphorus  in  air,  and  measuring  the 
residual  gas  (p.  416),  is  not  capable  of  application  in  an  exact  manner. 
It  is  better  to  use  a  large  amount  of  phosphorus  in  the  form  of  thin 
wire.  In  this  way  a  great  surface  is  obtained,  and  the  absorption  of 
oxygen  from  a  sample  of  air  may  be  carried  out  in  a  few  seconds. 
This  method  gives  fairly  accurate  results,  since  there  is  no  time  for  any 
appreciable  change  in  the  temperature  or  pressure  of  the  atmosphere 
during  the  experiment.  The  passage  of  purified  air  over  heated  copper 
(p.  416)  has  also  been  used  for  the  same  purpose.  When  this  method 
is  employed,  the  volume  of  the  nitrogen  and  argon  which  survives  the 
action  of  the  copper  is  measured,  while  the  increase  in  weight  of  the 
copper,  through  formation  of  cupric  oxide,  gives  the  weight  of  the 
oxygen  which  was  originally  mixed  with  it. 

Still  another  method,  which  is  in  constant  employment  in  the  analy- 
sis of  mixtures  of  gases,  may  also  be  applied  to  the  air.  It  consists  in 
mixing  a  measured  volume  of  air  with  an  excess  of  hydrogen,  reading 
off  the  volume  of  the  whole,  and  then  exploding  the  mixture  by  the 
passage  of  an  electric  spark.  A  tube  like  that  in  Fig.  44  (p.  125) 
may  be  used.  After  the  explosion,  the  steam  which  has  been  formed 
condenses.  The  contraction  gives  the  volume  of  the  gases  which  have 
disappeared  through  the  explosion,  and  of  this  one-third  is  oxygen  and 
two-thirds  hydrogen  (cf.  p.  125).  When,  for  example,  a  contraction 
of  25  c.c.  of  gas  has  occurred,  we  know  that  one-third  of  this  volume 
(namely,  8.3  c.c.)  is  the  volume  of  oxygen  which  the  measured  sample 
of  air  contained. 

In  the  air  taken  from  mines,  from  mountain  tops,  from  the  surface 
of  the  sea,  and  from  inland  regions,  the  proportion  of  oxygen  to  the 
residual  gas  is  found  to  be  fairly  constant,  although  easily  perceptible 
differences  are  noted.  The  percentage  of  oxygen  in  dried  air  ranges 
between  20.26  and  21.00,  the  latter  being  the  proportion  in  normal 
air. 

When  the  residual  gas  is  led  slowly  through  a  heated  tube  con- 


428  INORGANIC   CHEMISTRY 

taining  magnesium,  the  nitrogen  unites  with  the  metal  to  form  the 
solid  nitride  (p.  416),  and  only  about  10  c.c.  out  of  every  liter  remains 
uncombined.  This  residuum  is  argon,  mixed  with  one-hundredth  of 
its  volume  of  other  gases  belonging  to  the  same  family  (see  below). 
Exact  measurement  by  volume  gives  78.06  per  cent  of  nitrogen  and 
0.94  per  cent  of  argon  in  dried  air. 

It  is  possible  that  a  trace  of  hydrogen  (cf.  p.  92)  is  one  of  the 
regular  components  of  air. 

Components  wJiich  are  Variable  in  Amount.  —  Pure  country 
air  contains  about  3  parts  in  10,000  of  carbon  dioxide.  In  city  air 
there  are  from  6  to  7  parts  in  the  same  volume,  while  in  the  air  of 
audience-rooms,  where  the  ventilation  is  defective,  the  proportion  may 
rise  as  high  as  50  parts. 

The  simplest  way  of  showing  the  presence  of  carbon  dioxide  in  the 
air  is  by  exposing  a  solution  of  barium  hydroxide  in  a  shallow  vessel. 
After  a  short  time  a  layer  of  barium  carbonate  forms  upon  the  surface, 
in  consequence  of  a  chemical  change  represented  by  the  equation, 
Ba(OH)2  +  CO2  — >  BaC03  |  +  H20.  The  same  action  may  be  utilized 
for  the  purpose  of  quantitative  analysis.  A  measured  volume  of  air  is 
bubbled  slowly  through  a  measured  volume  of  a  solution  of  barium 
hydroxide  of  known  concentration,  and  the  quantity  of  barium  hy- 
droxide remaining  is  determined  by  titration  (p.  352). 

The  sources  of  the  carbon  dioxide  in  the  air  are  numerous.  It 
comes  from  the  decay  of  vegetable  and  animal  matter,  in  which, 
chiefly  through  the  influence  of  minute  vegetable  organisms,  the 
carbon  is  oxidized  to  carbon  dioxide.  It  is  formed  also  by  the  com- 
bustion of  coal  and  wood,  and  is  exhaled  by  animals.  The  proportion 
of  this  gas  in  the  air  would  naturally  increase  continuously,  though 
slowly,  as  the  result  of  these  processes,  were  it  not  that  it  is  removed 
just  as  continuously  by  the  action  of  growing  plants  (see  Chap,  xxviii). 
Thus,  there  is  no  reason  to  suppose  that  the  proportion  of  carbon 
dioxide  in  the  air  will  be  altered  appreciably,  the  two  kinds  of  opera- 
tions probably  balancing  one  another  approximately. 

The  quantity  of  water  vapor  in  the  air  is  constantly  changing.  It 
increases  locally  by  evaporation  from  the  soil  and  from  natural  waters, 
particularly  in  warm  weather.  It  decreases  when  local  cooling  leads 
to  the  precipitation  of  water  in  the  forms  of  mist  and  rain.  The 
phrase  commonly  heard,  that  on  a  moist  day  the  atmosphere  is  "  laden  " 
with  moisture,  is  peculiarly  inapt.  We  recognize  at  once  from  obser- 


THE   ATMOSPHERE.     THE   HELIUM   FAMILY  429 

vation  of  the  barometer,  which  is  lower  in  such  a  state  of  the  atmos- 
phere, that  the  pressure  of  the  air  is  less.  Moist  air  must  be  lighter 
than  dry  air,  for  in  it  a  certain  proportion  of  water  molecules,  of  molec- 
ular weight  18,  is  substituted  for  an  equal  number  (cf.  p.  190)  of  mole- 
cules of  nitrogen  and  oxygen  whose  relative  weights  are  28  and  32 
respectively.  The  result  is  therefore  a  diminution  in  the  specific 
gravity  of  the  air.  The  proportion  of  water  in  a  given  volume  of  air 
may  be  measured  most  accurately  by  permitting  the  air  to  stream 
slowly  through  tubes  filled  with  calcium  chloride  or  phosphoric  anhy- 
dride (Fig.  36,  p.  100).  The  increase  in  weight  of  the  charged  tubes 
represents  the  quantity  of  moisture  abstracted  from  the  sample. 

The  ammonium  nitrate  arises  from  the  interaction  of  nitric  acid 
and  ammonia.  The  latter  is  formed  by  the  decay  of  animal  matter 
(p.  417)  ;  the  former  by  the  union  of  nitrogen  and  oxygen  during 
thunder-storms.  The  electrical  discharges  produce  nitrogen  tetroxide 
(see  p.  439),  which  with  water  gives  nitric  acid  (q.v.). 

The  dust  varies  both  in  kind  and  quantity  according  to  the  locality. 
It  is  found  to  be  partly  inorganic.  In  the  country  the  nature  of  the 
inorganic  material  depends  upon  the  prevailing  components  of  the 
soil,  and  in  factories  the  dust  may  consist  of  minute  particles  of  glass, 
steel,  cement,  or  other  substances.  The  organic  dust  may  be  divided 
into  two  kinds.  The  part  which  is  dead  includes  coal  dust,  refuse 
from  the  streets,  minute  shreds  of  cotton,  linen,  hay,  etc.  The  living 
dust  consists  of  pollen  grains,  spores  of  fungi  and  other  plants,  germs 
of  infusoria,  and  bacteria.  The  presence  of  such  germs  in  the  air  is 
shown  by  the  fact  that  when  nutritive  liquids  have  been  exposed  to 
the  air,  even  for  a*few  minutes,  putrefaction  very  soon  sets  in.  Some 
of  these  germs  also  produce  disease  when  they  gain  access  to  the 
body,  particularly  through  wounds,  or  incisions  made  in  the  course  of 
operations.  The  object  of  antiseptic  treatment  by  the  use  of  phenol 
(carbolic  acid),  mercuric  chloride,  and  other  substances,  is  to  destroy 
such  organisms  or  to  hinder  their  development. 

Flasks  can  be  filled  with  dustless  air  through  the  displacement  of 
that  which  they  contain  by  air  drawn  through  a  wide  tube  packed  with 
12-15  inches  of  cotton.  It  has  been  shown  by  Aitken  that  air  of  this 
kind  behaves  differently  from  ordinary  air  in  respect  to  the  way  in 
which  its  moisture  condenses. 

If  a  sample  of  moist  air  is  cooled  until  it  contains  more  water  vapor 
than  it  could  take  up  at  the  existing  temperature,  the  excess  of  moist- 
ure is  deposited.  This  deposition  usually  takes  place  by  the  forma- 


430  INORGANIC   CHEMISTRY 

tion  of  a  multitude  of  little  particles  of  liquid  water,  which  together 
make  up  a  fog.  Now  dustless  air  lacks  this  property  entirely.  When 
saturated  with  water  and  then  cooled,  it  does  not  give  any  trace  of  fog. 
The  excess  of  moisture  is  gradually  deposited  upon  the  walls  of  the 
vessel  and  upon  any  material  objects  which  it  contains,  but  of  fog 
there  is  no  trace  visible.  It  seems  that  the  particles  of  dust  are  re- 
quired as  nuclei  round  which  the  water  may  gather.  In  the  absence 
of  dust,  and  therefore  of  proper  nuclei,  the  moisture  is  not  precipi- 
tated in  the  usual  way.  Thus  fogs  and  rain  would  be  impossible  but 
for  the  presence  of  dust  in  all  ordinary  air. 

By  diluting  air  with  dustless  air,  generating  fog  in  the  mixture, 
and,  with  the  help  of  a  microscope,  counting  the  globules  when  they 
settle,  an  estimate  of  the  number  of  particles  of  dust  in  air  may  be 
made.  It  is  found  that  rain  removes  a  large  proportion  of  them, 
while  respiration  and  combustion  greatly  increases  their  number.  The 
prevalence  of  fogs  in  cities  is  thus  accounted  for. 

NUMBER  OF 

DUST  PARTICLES 

IN  1  c.c. 

Outside,  raining 32,000 

Outside,  fair 130,000 

A  room 1,860,000 

A  room,  near  the  ceiling 5,420,000 

Air  above  Bunsen  flame 30,000,000 

Air  a  Mixture.  —  Since  the  main  components  of  air  were  not  defi- 
nitely identified  until  the  end  of  the  eighteenth  century,  we  can  under- 
stand why  the  substance  was  for  long  considered  to  be  an  element. 
The  experiments  which  we  have  described,  in  which  the  oxygen  was 
removed  from  the  air  and  the  nitrogen  remained,  do  not  prove  that 
the  original  constituents  were  present  simply  in  mechanical  mixture. 
They  might  have  been  combined,  and  the  combustion  of  phosphorus, 
for  example,  might  have  represented  the  removal  of  oxygen  from  com- 
bination with  nitrogen  and  its  appropriation  by  the  phosphorus.  It 
may  be  well,  therefore,  to  point  out  some  reasons  which  lead  us  to 
regard  the  air  as  a  mixture  : 

1.  When  oxygen  and  nitrogen  are  mixed  in  the  proper  proportions, 
we  obtain  a  gas  identical  with  air  in  all  its  properties,  and  there  is  no 
evidence  of  any  production  or  absorption  of  heat,  such  as  would  occur 
in  case  of  chemical  combination. 


THE   ATMOSPHERE.      THE   HELIUM   FAMILY  431 

2.  The  proportion  by  volume  in  which  the  gases  are  found  in  the 
air  is  not  so  simple  as  the  proportions  which  we  observe  in  cases  of 
chemical  combination.     The  proportion  is  close  to  4  :  1,  but  not  exactly 
4  :  1.    Besides,  as  we  have  seen,  the  proportion  is  not  perfectly  constant. 

3.  The  composition  of  air  varies,  while  the  composition  of  definite 
chemical  substances  is  always  the  same.     The  proportions  by  weight 
also  in  which  the  components  are  contained  in  air  are  not  integral 
multiples  of  the  atomic  weights. 

4.  When  two  substances  enter  into  chemical  combination,  the  new 
body  invariably  has  different  physical  properties  from  either  of  the 
original  ones.     Thus,  there  is  no  simple  relation  between  the  refrac- 
tive power  for  light  which  a  compound  possesses  and  the  refractive 
powers  of  its  constituents.     In  the  case  of  air,  however,  the  refrac- 
tive power  is  exactly  that  which  we  should  calculate  from  the  refrac- 
tive powers  of  the  constituents,  taking  into  account  the  proportions  of 
them  which  air  contains.     The  same  relation  holds  for  all  ordinary 
physical  properties  of  air.    For  example,  the  nitrogen  and  oxygen 
dissolve  independently  in  water  in  proportion  to  their  solubilities  and 
partial  pressures  (p.  155).     If  the  air  were  a  compound,  it  would  dis- 
solve as  a  whole,  and  the  relative  proportions  of  the  components  would 
not  be  changed  by  the  process. 

Composition  of  Air.  —  Air,  when  freed  from  carbon  dioxide  and 
water,  contains  by  volume  78.06  per  cent  of  nitrogen,  21.00  per  cent 
of  oxygen,  and  0.94  per  cent  of  argon.  When  only  the  water  is 
removed,  the  carbon  dioxide  averages  about  0.03  per  cent  of  the 
whole. 

Graham  has  suggested  an  illustration  which  will  make  those  pro- 
portions clearer.  He  says  that  if  we  imagine  the  air  to  be  divided  by 
magic  into  its  components,  and  to  remain  separated  for  a  time  suffi- 
ciently long  to  enable  us  to  note  the  proportions,  and  if  the  substances 
arrange  themselves  in  the  order  of  their  specific  gravities,  we  should 
have  the  following  layers  resting  upon  the  surface  of  the  earth  and 
upon  one  another  :  On  the  earth,  five  inches  of  water ;  above  that, 
thirteen  feet  of  carbon  dioxide ;  above  that,  a  mile  of  oxygen  ;  and  on 
the  top,  about  four  miles  of  nitrogen.  This  would  be  on  the  assump- 
tion that  these  gases  were  compressed  so  as  to  have  the  same  density 
throughout.  The  recent  discovery  of  argon  shows  that  we  should  add 
to  this  illustration  a  layer  of  argon,  of  about  ninety  yards  thickness, 
between  the  carbon  dioxide  and  oxygen. 


432  INORGANIC   CHEMISTRY 

Air  and  Health.  —  As  human  beings  we  have  an  especial  interest 
in  the  composition  of  the  air,  since  our  living  depends  upon  the  oxygen 
which  we  secure  by  breathing  it.  We  draw  about  half  a  liter  of  air 
into  our  lungs  at  each  breath,  or  about  half  a  cubic  meter  per  hour. 
The  oxygen  of  this  air  is  partly  used,  being  taken  up  by  the  blood, 
and  part  remains  in  the  exhaled  air.  On  the  other  hand,  carbon  dioxide 
is  given  off  in  the  lungs  and  passes  out  with  the  unused  oxygen.  The 
nitrogen  is  unaffected.  In  100  c.c.  of  expired  air  there  are  contained 
about  15.9  c.c.  of  oxygen  and  3.7  c.c.  of  carbon  dioxide.  The  total 
quantity  of  oxygen  consumed  during  twenty-four  hours  is  about  three- 
fourths  of  a  kilogram,  or  more  than  half  a  cubic  meter.  While  the 
greater  part  of  this  gains  access  to  the  body  through  the  lungs,  more 
or  less  exchange  of  gases  takes  place  in  all  animals  through  the  skin. 

The  lower  limit  of  oxygen  for  respirable  air  is  about  10  per  cent, 
although  a  candle  is  extinguished  if  the  proportion  falls  below  16.5  per 
cent.  The  union  of  oxygen  with  the  haemoglobin  of  the  red  blood- 
corpuscles,  and  of  the  carbon  dioxide  with  the  materials  in  the  plasma, 
are  reversible  reactions.  In  the  tissues,  where  the  concentration  of 
oxygen  is  small  and  that  of  carbon  dioxide  large,  the  former  leaves 
the  haemoglobin  and  the  latter  enters  the  plasma.  Conversely,  in 
the  lungs,  where  the  concentration  of  oxygen  is  at  a  maximum,  and 
that  of  carbon  dioxide  at  a  minimum,  the  reverse  changes  occur. 

Liquefaction  of  Gases.  —  The  earliest  experiments  of  this  kind 
seem  to  have  been  made  by  Northmore  (1805),  who  liquefied  chlorine, 
hydrogen  chloride,  and  sulphur  dioxide.  In  1823  chlorine  was  again 
liquefied  by  Faraday;  and  in  the  same  year  Davy,  whose  assistant 
Faraday  was,  liquefied  hydrogen  chloride.  During  the  following  years 
Faraday  reduced  other  gases  —  sulphur  dioxide,  hydrogen  sulphide, 
carbon  dioxide,  nitrous  oxide,  cyanogen,  and  ammonia — to  the  liquid 
condition. 

The  method  which  he  employed  was  extremely  simple.  He  used  a 
bent  tube  shaped  like  an  inverted  v  (A),  into  one  limb  of  which  materials 
for  producing  the  gas  were  introduced.  The  other  limb  was  then 
sealed  up  and  immersed  in  a  freezing  mixture.  The  gas,  usually 
liberated  by  heating,  was  liquefied  by  its  own  pressure  in  the  cold 
limb.  By  means  of  a  more  elaborate  apparatus,  Cailletet  and  Pictet 
simultaneously  (December,  1877)  obtained,  the  one,  a  fog,  and  the 
other,  a  spray  containing  droplets  of  liquid  oxygen.  In  1883  Wrob- 
lewski  and  Olszewski  made  larger  amounts  of  the  same  liquid.  About 


THE   ATMOSPHERE.     THE    HELIUM   FAMILY 


433 


the  same  time  Dewar  devised  means  of  manufacturing  large  quantities 
of  liquid  air  and  oxygen. 

The  principle  now  used  in  liquefying  gases  depends  on  the  fact 
that  a  perfect  gas,  when  expanding  into  a  vacuum,  should  suffer  no 
fall  in  temperature,  since  it  does  no  work,  while  ordinary  gases 
do  become  cooled  very  slightly.  The  work  which  they  do  in  expand- 
ing in  such  circumstances  is  done  in  over- 
coming the  cohesion  between  their  mole- 
cules (p.  131),  and  occasion  for  it  arises 
from  the  fact  that  a  tearing  apart  of  the 
substance,  which  consumes  heat,,  has  to 
take  place.  Since  this  cohesion  becomes 
more  conspicuous  the  lower  the  tempera- 
ture (cf.  p.  135),  the  cooling  effect  of  ex- 
pansion becomes  greater  and  greater  as  the 
temperature  falls. 

By  ingenious  arrangement  of  the  appa- 
ratus, the  low  temperature  produced  by  the 
expansion  of  the  compressed  gas  is  used  to 
cool  fresh  portions  of  the  gas  which  have 
not  yet  been  expanded.  The  tube  contain- 
ing the  compressed  gas  is  placed  concen- 
trically within  another  tube,  and  the 
stream  of  gas  after  it  has  expanded  is 
turned  back  upon  its  course  and  passes 
through  the  annular  space  along  the  tube 
containing  the  unexpanded  gas.  This 
double  tube  is  many  yards  long,  so  that  the 

most  perfect  contact  and  complete  interchange  of  heat  can  occur.  Thus, 
by  a  process  of  intensification,  the  temperature  of  the  unexpanded  gas 
is  continually  depressed,  and  the  temperature  after  expansion,  there- 
fore, becomes  lower  still.  Finally,  the  latter  temperature  becomes 
low  enough  to  produce  liquefaction  of  the  compressed  gas  at  the 
pressure  to  which  it  is  subjected,  and  the  jet  of  expanding  gas  is  filled 
with  drops  of  liquid.  These  are  caught  in  a  recess  of  the  vessel,  and 
may  be  removed  from  time  to  time  through  a  stopcock. 

The  most  successful  apparatus  for  use  on  a  small  scale  is  that  de- 
vised by  Hampson  (Fig.  86).  The  compressed  gas  enters  by  the  small 
tube  near  the  top,  and  eventually  issues  by  the  wider  one.  The  inner 
tube  is  coiled  in  a  small  cylinder  to  secure  compactness,  and  a  spiral 


FIG.  86. 


434 


INORGANIC   CHEMISTRY 


partition  between  the  coils  produces  the  outer  tube  of  which  we  have 
spoken.  The  gas  in  the  tube  A  (Fig.  87)  is  under  a  pressure  of  150- 
200  atmospheres.  The  distance  of  the  nozzle  D  from  the  plug  C  is 
adjusted  so  that  the  pressure  of  the  gas  in  the 
chamber  and  spiral  outer  tube  is  reduced  to  one 
atmosphere. 

For  handling  liquefied  gases,  a  form  of  appara- 
tus devised  independently  by  Weinhold  and  Dewar 
is  employed  (Fig.  88).  It  consists  of  a  double 
flask  in  which  the  space  between  the  inner  and 
outer  bulbs  has  been  exhausted  by  means  of  an 
air-pump.  To  prevent  by  reflection  the  access  of 
radiant  heat  to  the  interior,  the  surfaces  of  the 
flasks  are  often  silvered. 


FIG.  87. 


Liquid  Air.  —  Liquid  air  varies  in  composition, 
as  the  nitrogen  (b.-p.  —  194°)  is  less  condensible 
than  the  oxygen  (b.-p.  —  182.5°).  It  boils  at  about 
—  190°,  and  contains  about  54  per  cent  of  oxygen 
by  weight,  while  air  contains  23.2  per  cent.  By 
allowing  evaporation  to  go  on,  a  liquid  containing 
75  to  95  per  cent  of  oxygen  is  easily  obtained  (cf. 
p.  63).  The  gas  secured  by 
the  evaporation  of  the  residue 

is  pumped  into  cylinders  and  sold  as  compressed 

oxygen.      Cartridges   made  of  granular  charcoal 

and  cotton  waste,  when  saturated  with  liquid  air, 

have  been  used  as  an  explosive  in  mining. 

THE   HELIUM   FAMILY. 

Argon.  —  Lord  Rayleigh  was  the  first  to  ob- 
serve that,  while  specimens  of  oxygen  and  other 
gases  made  purposely  from  various  sources  always 
had  the  same  density,  nitrogen  was  an  exception. 
One  liter  of  nitrogen  made  from  air,  and  supposed  to  be  pure,  weighed 
1.2572  g.  When  the  gas  was  manufactured  by  decomposition  of  five 
different  compounds,  such  as  urea  and  certain  oxides  of  nitrogen,  the 
results  agreed  well  amongst  themselves.  The  mean  weight  of  a  liter 
of  this  nitrogen  was  only  1.2505  g.  The  difference,  amounting  to  nearly 


FlG. 


THE   ATMOSPHERE.     THE  HELIUM  FAMILY  435 

7  nig.,  was  very  much  greater  than  the  experimental  error.  The  sus- 
picion naturally  arose  that  some  heavier  gas  was  present  in  natural 
nitrogen.  Soon  after  (1894),  Professor,  now  Sir  William  Ramsay 
obtained  argon  by  removal  of  the  greatly  preponderating  nitrogen  by 
means  of  magnesium  (p.  427).  The  new  gas  had  a  molecular  weight  of 
about  40,  and  was  therefore  more  than  one-third  heavier  than  nitrogen. 

In  order  to  make  sure  that  this  substance  did  not  have  its  source  in  the  magne- 
sium, a  different  method  was  used  by  Lord  Rayleigh  to  separate  it  from  nitrogen. 
He  inclosed  the  nitrogen  with  a  sufficient  quantity  of  oxygen  in  a  flask,  through 
the  sides  of  which  platinum  poles  had  been  inserted.  A  tube  entered  the  flask  by 
the  neck,  and  through  this  a  constant  fountain  of  potassium  hydroxide  solution 
played  upon  the  interior  and  kept  the  surface  covered  with  fresh  quantities  of  the 
liquid.  Another  tube  permitted  the  overflow  of  the  excess  of  this  solution.  The 
discharge  of  electricity  produced  nitrogen  tetroxide  (q.v.),  which  was  absorbed  by 
the  potassium  hydroxide  to  form  potassium  nitrate  and  potassium  nitrite.  The 
volume  of  gas  thus  continually  diminished,  and,  by  persistent  sparking  of  the  mix- 
ture with  oxygen,  the  nitrogen  was  finally  all  taken  out.  The  excess  of  oxygen 
was  then  removed,  and  the  gas  which  remained  was  found  to  be  identical  with  that 
which  Kamsay  had  obtained. 

Lord  Rayleigh's  method  was  extremely  interesting,  since  it  was  a  reproduction 
of  an  experiment  made  by  Cavendish  towards  the  end  of  the  eighteenth  century.. 
The  latter  had  remarked  that  the  assumption  that  the  inert  atmospheric  gas  was  a 
homogeneous  single  substance  had  not  been  confirmed  by  sufficiently  careful  ex- 
periment. He  even  endeavored  in  precisely  the  above  way  to  remove  the  nitro- 
gen in  order  to  see  whether  any  other  body  remained.  He  records  the  fact  that 
the  residual  gas  was  nothing  but  a  minute  bubble,  and  seems  to  have  dismissed  the 
subject  with  the  idea  that  if  there  was  any  other  constituent  of  the  atmosphere 
present  in  the  nitrogen,  its  amount  was  exceedingly  small.  Argon  thus  narrowly 
escaped  detection  nearly  a  century  before  its  actual  discovery. 

The  exact  density  of  argon,  referred  to  oxygen  =  32,  is  39.9. 
When  liquefied  it  boils  at  —  186°,  and  the  colorless  solid,  obtained  by 
cooling  the  liquid,  melts  at  —  189.5°.  The  solubility  of  the  gas  in 
water  (4  volumes  in  100)  is  two  and  one-half  times  that  of  nitrogen.  It 
has  not  been  found  to  enter  into  any  sort  of  chemical  combination,  and 
was  named  argon  on  this  account  (Gk.  a/ryo?,  inactive). 

Since  the  atomic  weight  of  a  substance  is  a  quantity  showing  the 
proportion  in  which  it  enters  into  combination,  it  will  be  seen  that 
argon,  since  it  has  not  yet  been  found  to  combine  with  anything,  has, 
to  speak  strictly,  no  atomic  weight  (pp.  200,  205).  At  the  same  time,  it 
is  manifestly  a  question  of  interest  to  determine  whether  the  physical 
properties  require  the  supposition  that  the  molecule  of  argon  contains 
one  atom,  or  more  than  one  atom. 

If  gases  were  composed  of  perfectly  elastic  spheres,  the  molecules 


436  INORGANIC   CHEMISTRY 

would  be  altered  only  in  respect  to  velocity  of  movement  by  heating. 
Calculation  enables  us  to  determine  that  to  raise  the  temperature  of 
one  G.M.V.  of  such  a  gas  by  one  degree  would  require  3  calories  in 
every  case.  Now  Regnault  found  the  following  values  (in  calories)  for 
the  heat  capacity  of  gases  : 

Oxygen  (02) 4.96  Carbon  dioxide  (C02)  .     7.56 

Hydrogen  (H,)     ....  4.82  Sulphur  dioxide  (S02)?      7.82 

Nitrogen  (N2)      ....  4.82  Chloroform  (CHC18)      .  16.55 

Hydrogen  chloride  (HC1),    4.76  Alcohol  (C2H60)      .     .  18.70 

These  gases  evidently  are  not  constituted  as  the  hypothesis  with 
which  this  paragraph  opened  supposes.  Some  heat  is  consumed  in 
work  done  inside  the  polyatomic  molecules,  and  the  amounts  by  which 
the  numbers  exceed  3  calories  show  that  the  intramolecular  work  is 
greater  as  the  complexity  of  the  molecules  increases.  Now,  in  mercury 
vapor  the  value  is  exactly  3,  and  we  have  already  seen  that  its  atomic 
and  molecular  weights  are  identical  (p.  204).  According  to  the  molec- 
ular hypothesis,  its  molecules  are  monatomic,  and  in  them  there  is  no 
opportunity  for  the  consumption  of  heat  in  intramolecular  change. 
Hence,  when  argon  was  found  likewise  to  give  3  for  the  value  of  its 
molecular  heat-capacity,  identity  of  its  atomic  and  molecular  weights 
was  assumed  also. 

Helium.  —  In  1868  Lockyer  first  detected  an  orange  line  in  the 
spectrum  of  the  sun's  prominences  which  was  not  given  by  any  terrestrial 
substance  then  known.  The  line  was  so  conspicuous  that  it  was  attrib- 
uted to  the  presence,  in  considerable  quantity,  of  a  new  chemical  ele- 
ment, which  was  named  helium  (Gk.  ^Atos,  the  sun).  Ramsay,  in 
searching  for  sources  of  argon,  examined  the  "  nitrogen  "  which  was  re- 
ported by  various  mineralogists  as  being  disengaged  when  certain  rare 
minerals  were  heated.  These  minerals,  cleveite,  uraninite,  and  brog- 
gerite,  were  chiefly  compounds  of  uranium,  yttrium,  and  thorium.  He 
was  surprised  to  find  (1895)  that  the  gas  was  not  always  nitrogen,  nor 
was  it  even  argon.  It  frequently  contained  a  large  proportion  of  a  gas, 
very  much  lighter  than  either,  the  spectrum  of  which  showed  at  once  that 
it  was  identical  with  helium.  The  same  gas  has  since  been  obtained 
from  the  water  of  certain  mineral  springs,  and  is  found  in  small  amount 
in  the  atmosphere.  Helium  does  not  exhibit  any  tendency  to  enter  into 
combination,  either  with  the  elements  which  its  parent  minerals  con- 
tain, or  with  any  others.  It  is  monatomic  (cf.  p.  435)  ;  its  density  shows 
that  its  molecular  weight  is  4.  It  has  been  liquefied  (b.-p.  4.5°  Abs.)0 


THE  ATMOSPHERE.   THE  HELIUM  FAMILY        437 

Neon,  Krypton,  and  Xenon.  —  When  the  argon  obtained  from 
atmospheric  nitrogen  is  cooled  with  liquid  air  (  —  185°),  the  argon, 
krypton,  and  xenon  are  liquefied,  and  the  neon  and  helium  are  dissolved 
by  the  liquid.  When  heat  is  allowed  to  reach  the  mixture,  the  last 
two  gases  escape  first,  along  with  much  argon.  When  most  of  the 
argon  has  escaped,  the  krypton  and  xenon  still  remain  liquid.  By 
repeated  liquefaction  and  fractional  evaporation  (see  under  Petroleum), 
the  krypton  and  xenon  are  separated  from  the  argon  and  from  one 
another.  When  the  vessel  containing  the  mixture  of  helium  and  neon 
is  immersed  in  liquid  hydrogen  (—240°),  the  second  freezes  to  a  white 
solid,  and  the  helium,  which  remains  gaseous,  can  be  pumped  off. 

These  gases  are  all  entirely  inactive  chemically,  and  are  all  mon- 
atomic.  Their  molecular  weights  are  :  Neon,  20;  krypton,  81.5;  xenon, 
128. 

Exercises.  — 1.  A  sample  of  moist  air,  confined  over  water  at  15° 
and  760  mm.,  occupies  15  c.c.  It  is  mixed  with  20  c.c,  of  hydrogen,  and 
the  mixture  is  exploded,  and  suffers  a  contraction  of  9.5  c.c.  What 
would  be  the  volume  of  the  oxygen  it  contained  if  measured  dry  at  0° 
and  760  mm.  ? 

2.  Calculate  from  the  data  on  pp.  427,  428  and  the  densities  the 
percentage  by  weight  of  the  three  principal  components  of  air. 

3.  Which  of  the  proofs  that  air  is  a  mixture  (p.  430)  show  simply 
that  the  components  are  not  wholly  combined,  and  which  show  that 
not  even  a  part  is  combined  ? 


CHAPTER   XXVI 
OXIDES   AND    OXYGEN  ACIDS   OF   NITROGEN 

THE  names  and  formulae  of  the  oxides  and  oxygen  acids  of  nitrogen 
are  as  follows : 

Nitrous  oxide  N20  < Hyponitrous  acid  H^N^ 

Nitric  oxide  NO 

Nitrous  anhydride  N80g  < >      Nitrous  acid  HNO2 

Nitrogen  tetroxide  N204  and  N02 

Nitric  anhydride  N2O6     < —       — >      Nitric  acid  HNO8. 

All  the  oxides  are  endothermal  compounds,  yet,  with  the  exceptions  of 
the  third  and  the  last,  they  are  all  relatively  stable.  The  acids,  when 
deprived  of  the  elements  of  water,  yield  the  oxides  opposite  which 
they  stand.  Conversely,  excepting  in  the  case  of  nitrous  oxide,  the 
anhydrides  with  water  give  the  acids.  All  of  these  substances  are 
obtained  directly  or  indirectly  from  nitric  acid  —  nitric  anhydride  by 
removal  of  water,  the  others  by  reduction.  We  turn,  therefore,  first, 
to  this  acid,  its  sources  and  properties. 

NITRIC  ACID. 

Sources.  —  Sodium  nitrate,  or  Chili  saltpeter,  is  found  in  a  des- 
ert region  near  the  boundary  of  Chili  and  Peru,  and  chiefly  in  the 
former  country.  The  deposit  is  about  5  feet  thick,  2  miles  wide, 
220  miles  in  length,  and  contains  20  to  56  per  cent  of  the  salt.  Puri- 
fication is  effected  by  recrystallization.  Potassium  nitrate,  or  Bengal 
saltpeter,  is  found  in  the  soil  in  the  neighborhood  of  cities  in  India, 
Persia,  and  other  oriental  countries.  It  arises  from  the  oxidation  of 
animal  refuse  (cf.  p.  417)  through  the  mediation  of  nitrifying  bacteria. 
The  potash  and  lime  in  the  soil,  along  with  the  product  of  oxidation  of 
the  nitrogen,  give  nitrates  of  potassium  and  calcium.  The  aqueous 
extract  of  this  soil  is  treated  with  wood  ashes,  on  account  of  the  potash 
(K2C08)  contained  in  them,  is  poured  'off  from  the  calcium  carbonate 
thus  precipitated,  and  is  finally  evaporated. 

438 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  439 

The  action  of  the  nitrifying  bacteria  may  bo  imitated  in  a  rough  way.  Air  is 
caused  to  pass  slowly  through  concentrated  aqueous  ammonia,  whereby  it  becomes 
mixed  with  ammonia  gas.  This  mixture  is  led  through  a  wide  tube  containing 
platinized  asbestos  and  is  then  discharged  into  a  large  flask.  When  the  asbestos  is 
warmed,  it  begins  to  glow,  and  thereafter  the  action  maintains  itself.  A  part  of 
the  ammonia  is  oxidized  to  nitric  acid,  which  combines  with  the  excess  of  ammonia, 
giving  ammonium  nitrate.  This  salt  forms  a  cloud  which  settles  in  solid  form  in 
the  flask. 

Preparation.  —  When  any  nitrate  is  treated  with  any  acid,  nitric 
acid  is  formed  by  a  reversible  double  decomposition.  As  sodium 
nitrate  is  the  cheapest  salt  of  nitric  acid,  it  is  always  employed.  For 
the  same  reason,  and  on  account  of  its  activity,  and,  above  all,  because 
of  its  relative  involatility,  sulphuric  acid  is  used  to  displace  it  : 

NaN08  +  ttjSO,  *=?  NaHS04  +  HN08f  . 

The  nitric  acid  is  rather  volatile  (b.-p.  86°),  while  sulphuric  acid  (b.-p. 
330°)  is  much  less  so,  and  the  two  salts  are  not  volatile  at  all.  Thus 
the  interaction  proceeds  to  completion  very  easily  (cf.  p.  259).  The 
materials  are  heated  in  cast-iron  stills,  and  the  vapor  is  condensed  in 
earthenware  pipes  surrounded  by  water.  In  many  factories  a  reduced 
pressure  is  maintained  in  the  stills  and  condensers,  in  order  that  the 
distillation  may  take  place  at  the  lowest  possible  temperature.  This 
precaution  is  taken  to  reduce  to  a  minimum  the  partial  decomposition 
of  the  nitric  acid  (see  below). 

Another  action  by  which  attempts  are  being  made  to  manufacture 
nitric  acid  is  the  direct  union  of  the  nitrogen  and  oxygen  of  the  air 
under  the  influence  of  an  electric  discharge.  The  nitrogen  tetroxide 
(K02),  which  is  formed  in  small  amounts  at  a  time,  is  dissolved  in 
water  : 

3N02  +  Hp  ?±  2HN08 


The  nitric  oxide  gas,  on  escaping  from  the  water,  unites  directly  with 
oxygen  to  reproduce  the  tetroxide.  The  reaction  is  of  interest,  inde- 
pendently of  this  one  application,  because  of  its  reversibility.  It  pro- 
ceeds forward  with  excess  of  water,  while  the  reverse  action  takes 
place  when  nitric  oxide  (y.v.)  comes  in  contact  with  concentrated  nitric 
acid  in  which  the  quantity  of  water  is  at  a  minimum. 

Physical  Properties.  —  Nitric  acid  is  a  colorless,  mobile  liquid, 
boiling  at  86°,  and  freezing  to  a  solid  which  melts  at  —  47°.     It  fumes 


440  INORGANIC   CHEMISTRY 

strongly  when  its  vapor  issues  into  moist  air  (of.  p.  182).  An  aqueous 
solution  containing  68  per  cent  of  the  acid  boils  at  120.5°,  while  the 
pure  acid,  pure  water,  and  all  other  mixtures,  boil  at  lower  temperatures, 
and  have,  therefore,  higher  vapor  pressures.  On  this  account  a  more 
dilute  acid,  when  heated,  loses  water  until  it  reaches  this  strength 
(of.  p.  182).  The  68  per  cent  nitric  acid  forms  the  "  concentrated 
nitric  acid  "  of  commerce. 

Chemical  Properties.  —  1.  Like  chloric  acid  (p.  274),  and  other 
oxygen  acids  of  the  halogens,  nitric  acid  is  most  stable  when  mixed 
with  water.  The  pure  (100  per  cent)  acid  decomposes  while  being 
distilled  : 

4HN03  ->  4N02  -f  2H20  +  02, 

yet  not  with  explosive  violence  like  chloric  acid.  The  distillate  is 
colored  brown  by  dissolved  nitrogen  tetroxide  (N02).  Repeated  dis- 
tillation finally  leaves  68  per  cent  of  the  acid,  mixed  with  32  per  cent 
of  water  formed  by  the  above  decomposition.  The  acid  of  constant 
boiling-point  is,  therefore,  reached,  as  usual,  from  more  concentrated  as 
well  as  from  less  concentrated  specimens. 

"  Fuming"  nitric  acid  is  brown  in  color,  and  contains  a  considerable 
amount  of  dissolved  nitrogen  tetroxide.  It  is  made  by  distilling  the 
acid  with  a  little  starch.  The  latter  reduces  a  part  of  the  nitric  acid 
and  liberates  more  of  the  tetroxide  than  does  mere  distillation. 

2.  Nitric  acid  combines  with  small  amounts  of  water  to   form 
hydrates  HN08,H20  and  HN08,3H20,  but  they  are  unstable,  and  are 
decomposed  when  more  water  is  added.     The  formula  of  the  former 
might  be  written  H8N04,  but  no  salts  corresponding  to  a  tribasic  acid 
of  this  constitution  are  known  (see  Phosphoric  acid). 

3.  Nitric  acid,  when  dissolved  in  water,  is  highly  ionized,  and  gives 
a  solution  containing  a  relatively  large  concentration  of  hydrion.     It  is 
therefore  active  as  an  acid.     By  interaction  with  hydroxides  and  oxides 
it  forms  nitrates. 

4.  When  pure  nitric  acid  (b.-p.  86°)  is  poured  upon  phosphoric 
anhydride,  the  latter  possesses  itself   of  the  elements  of  water,  and 
distillation  of  the  mixture  gives  nitric  anhydride  : 


2HN08  -f  PA  ->  N206  1  +  2HP08. 

The  anhydride  is  a  white  solid  melting  at  30°  and  boiling  at  45°.     It 
unites  vigorously  with  water  to  form  nitric  acid.     It  cannot  be  kept,  as 


OXIDES   AND   OXYGEN   ACIDS   OF   NITROGEN  441 

it  decomposes  into  nitrogen  tetroxide  and  oxygen,  2N205  — •»  4NO2  -f-  Oa, 
with  liberation  of  heat. 

5.  Like  the  unstable  oxygen  acids  of  the  halogens,  nitric  acid  ifi  an 
oxidizing  agent  even  when  diluted  with  water.     The  multiplicity  of 
the  products  into  which  it  may  be  decomposed  by  reduction,  however, 
renders  separate  treatment  of  this  property  necessary  (see  below). 

6.  Nitric   acid  interacts    energetically  with    many  compounds  of 
carbon.     Thus,  when  heated  with  phenol  (carbolic  acid)  it  gives  picric 
acid,  which  crystallizes  in  yellow  needles  in  the  mixture  : 

C6H6(OH)  +  3HON02  -*  C6H2(OH)(N02)3  +  3H2O. 

The  presence  of  water  decreases  the  activity  of  the  molecules.  Hence, 
in,  this  sort  of  action,  which  is  not  ionic,  not  only  is  the  most  concen- 
trated nitric  acid  employed,  but  concentrated  sulphuric  acid  is  added  to 
assist  in  the  elimination  of  the  water  (cf.  p.  388)  that  arises  as  one 
of  the  products. 

It  will  be  seen  that  the  group  N02  has  taken  the  place  of  hydro- 
gen which  was  formerly  attached  directly  to  the  carbon  of  the  phenol. 
Compounds  of  this  kind  are  called  nitro-derivatives.  Picric  acid  is 
trinitrophenol. 

7.  Organic  compounds  of  another  class,  the  alcohols  (q.v.),  interact 
with  molecular  nitric  acid  in  a  different  way.     The  latter  is  mixed  with 
sulphuric  acid  with  the  same  object  as  before.     Thus,  when  glycerine 
is  added  slowly  to  the  cooled  mixture,  glyceryl  nitrate  (so-called  nitro- 
glycerine, see  below)  is  produced  : 

C3H5(OH)8  +  3HONO2  ->  C3H5(ON02)3  -f  3H20. 

Here  it  is  the  hydrogen  of  the  hydroxyl  groups  that  is  displaced  by 
N02.  The  action  is  not  ionic,  and  the  product  is  not  an  ionogen. 
Gun-cotton  is  made  by  this  action,  prepared  cotton  (cellulose)  being 
employed  : 

(C6H,A),  +  6HONO,  '-4  C12HU0IO(N02)8  +  6H2O. 

8.  Nitric  acid  produces  substances  of  bright-yellow  color,  known  as 
xanthoproteic  acids,  when  it  comes  in  contact  with  the  skin. 

The  chemical  properties  of  nitric  acid  are  best  represented  by  the 
graphic  formula : 


442  INORGANIC   CHEMISTRY 

Nitrates.  —  The  nitrates  are  all  more  or  less  easily  soluble  in 
water.  When  heated  they  decompose  (see  below).  Sodium  nitrate  is 
used  most  largely  as  a  fertilizer.  Much  is  employed  in  sulphuric  acid 
manufacture,  and  the  rest  for  conversion  into  potassium  nitrate  and  in 
making  nitric  acid.  Potassium  nitrate  is  used,  along  with  sulphur  and 
charcoal,  in  the  manufacture  of  gunpowder.  It  furnishes  the  oxygen 
with  which  the  charcoal  unites  ;  and  potassium  sulphide,  carbon  dioxide, 
and  nitrogen  are  amongst  the  products  of  the  explosion. 

NITRIC  OXIDE  AND  NITROGEN  TETROXIDE. 

Preparation  of  Nitric  Oxide.  —  Pure  nitric  oxide  is  obtained 
by  adding  nitric  acid  to  a  boiling  solution  of  ferrous  sulphate  in  dilute 
sulphuric  acid  or  of  ferrous  chloride  in  hydrochloric  acid  : 

2FeS04  +  H2S04  ->  Fe2(S04)3  ( +  2H)   x  3  (1) 

(3H)  +  HN03  ->  NO  +  2H2O  x  2  (2) 

6FeSO4  +  3H2SO4  +  2HN08  ->  3Fe2(S04)8  +  2NO~+~lH2OT 

The  first  partial  equation  does  not  take  place  at  all  unless  an  oxidizing 
agent  like  nitric  acid  is  present  (p.  307).  The  multiplication  of  the 
two  partial  equations  by  3  and  2  respectively  is  required  in  order  that 
the  hydrogen,  which  is  not  a  product,  may  cancel  out.  This  action  is 
used  as  a  means  of  determining  the  quantity  of  nitric  acid  in  a  solution, 
or  of  nitrates  in  a  mixture,  by  measurement  of  the  volume  of  nitric 
oxide  evolved. 

As  we  shall  see,  this  gas  may  also  be  obtained  when  sufficiently 
dilute  nitric  acid  (sp.  gr.  1.2)  acts  upon  copper.  Although  some 
nitrous  oxide  and  nitrogen  are  produced  in  this  interaction,  it  fur- 
nishes a  convenient  method  of  generating  the  gas. 

Properties  of  Nitric  Oxide.  —  Nitric  oxide  is  a  colorless  gas.  In 
solid  form  it  melts  at  — 150°  and  the  liquid  boils  at  — 142.4°  under  757.2 
mm.  pressure.  Its  solubility  in  water  is  slight. 

The  density  of  the  gas  shows  the  formula  to  be  NO  ;  and  there  is  no 
tendency  to  form  a  polymer,  such  as  N202,  even  at  low  temperatures. 
This  gas  decomposes  when  heated,  until  only  the  proportion  in 
equilibrium  at  the  existing  temperature  remains.  The  action  is 
reversible.  When  air  (containing  excess  of  nitrogen,  so  far  as  this 
action  is  concerned)  is  heated,  the  proportions  of  nitric  oxide 
formed  are :  at  1922°,  one  per  cent,  and  at  2927°  about  five  per  cent. 


OXIDES  AND   OXYGEN   ACIDS   OF   NITROGEN  443 

Vigorously  burning  phosphorus  continues  to  burn  in  the  gas,  the 
heat  evolved  liberating  the  oxygen  required  for  the  continuation  of  the 
combustion.  Burning  sulphur  and  an  ignited  taper,  however,  are 
extinguished. 

Nitric  oxide  has  two  characteristic  properties.  It  unites  directly 
with  oxygen  in  the  cold  to  form  the  reddish-brown  nitrogen  tetroxide  : 

2NO  +  02  <=±  2N02. 

The  same  result  follows  when  it  is  led  into  warm  concentrated  nitric 
acid  (cf.  p.  439)  : 

NO  +  2HNOS  <=±  3N02  +  H2O. 

It  also  unites  with  a  number  of  salts,  the  compound  in  the  case  of 
ferrous  sulphate  being  capable  of  existence  in.  solution  and  possessing 
a  brown  color.     The  composition  of  this  compound  (a  molec- 
ular compound,  see  below)  has  not  been  determined,  but  at 
8°  the  proportion  of  nitric  oxide  absorbed  by  the  solution  is 
about  2NO  :  3FeS04. 

Since  ferrous  sulphate  will  first  reduce  nitric  acid  to  nitric 
oxide  (p.  442),  and  the  excess  of  the  salt  will  then  give  a 
brown  color  with  the  product,  a  delicate  test  for  nitric  acid  is 
founded  upon  the  above  action.  The  substance  supposed  to 
contain  a  nitrate  is  mixed  with  a  strong  solution  of  ferrous 
sulphate,  and  concentrated  sulphuric  acid  is  poured  down  the 
side  of  the  tube  so  as  to  lie  below  the  lighter  mixture  (Fig. 
89).  At  the  surface  of  contact  the  sulphuric  acid  liberates 
the  nitric  acid,  and  a  brown  layer  is  seen.  Even  when  the 
amount  of  the  nitrate  is  very  small,  the  brown  tint  may  be 
distinctly  made  out  by  contrast  with  the  nearly  colorless 
liquids  above  and  below  it.  FIG.  89. 

Molecular  Compounds.  —  When  substances  formed  by  union  of 
two  compounds  have  a  prevailing  tendency  to  decompose  into  the 
same  two  materials,  and  exhibit  the  chemical  properties  of  their  con- 
stituents rather  than  individual  ones  of  their  own,  they  are  often 
called  molecular  compounds.  Thus  the  above  substance,  3FeS04> 
2NO,  gives  off  the  nitric  oxide  again  when  warmed,  and  its  solution 
has  the  properties  of  a  mixture  of  ferrous  sulphate  and  nitric  oxide. 
Similarly,  hydrates  (pp.  120, 123)  are  formed  by  union  of  salts  or  other 
substances  with  water,  and  are  apparently,  for  the  most  part,  decora- 


444  INORGANIC   CHEMISTRY 

posed  by  solution.  Double  salts  (p.  360),  such  as  ferrous-ammonium 
sulphate  FeS04,(NH4)2S04,  6H30,  of  which  very  many  are  known,  are 
of  the  same  character.  They  are  stable  only  in  the  solid  form.  There 
are  also  compounds  of  salts  with  ammonia  (see  Compounds  of  copper 
and  silver),  and  with  carbon  monoxide  (CO),  one  such  compound  being 
formed  with  cuprous  chloride  (q-v.). 

The  name  molecular  compounds  is  derived  from  the  supposition 
that,  in  these  compounds,  the  molecules  of  the  components  retain  their 
integrity  to  some  extent  and  are  thus  ready  to  be  liberated.  This  is 
an  attempt  to  explain  the  fact  that  the  behavior  is  that  of  the  con- 
stituents. It  distinguishes  molecular  compounds  from  substances  like 
ammonium  chloride  and  phosphorus  pentachloride.  The  former  may 
be  made  by  union  of  HC1  and  NH3,  but  usually  behaves  rather  as  if 
composed  of  NH4  and  Cl.  The  latter  (q.v.)  dissociates  into  PC18  and 
Cla,  but  with  water  gives  phosphoric  acid  (cf.  p.  181),  which  is  derivable 
from  the  pentachloride  only.  The  distinction  is  of  practical  rather 
than  theoretical  importance,  however,  for  there  are  all  gradations  in 
the  behavior  of  molecular  compounds.  It  is  useful  simply  as  a  rough 
means  of  classifying  and  remembering  certain  facts. 

Distinguishing  molecular  compounds  from  ordinary  compounds  is 
further  justified  by  the  fact  that  the  constituents  of  molecular  com- 
pounds often  seem  to  be  saturated  (p.  379),  and  no  ordinary  valences 
are  available  for  holding  the  new  material.  Thus  in  Ca11^1  the 
ordinary  valences  are  all  saturated.  Yet  the  salt  forms  the  hydrate 
CaCl2,6H20  with  water  (H^O11)  which  is  likewise  a  saturated  compound. 
The  conception  of  molecular  compounds  implies,  therefore,  the  idea  of 
a  sort  of  valence  of  molecules.  Thus  FeSO4  forms  FeS04,7H20  and 
FeS04,(NH4)2S04,6H20,  and  FeS04,K2S04,6H20,  in  all  of  which 
sepen  other  molecules  are  combined  with  it.  The  sulphates  of  other 
bivalent  metals  (q.v.)9  such  as  copper  and  magnesium,  form  molecular 
compounds  of  the  same  nature.  Ammonium  chloride,  on  the  other 
hand,  is  not  a  molecular  compound,  because,  although  NH8  unites  with 
HC1,  HBr,  HI,  and  HE,  yet  nitrogen  is  quinquivalent,  and  substances 
like  N205,  NH4C1,  etc.,  may  fitly  be  regarded  as  ordinary  compounds. 

Preparation  of  Nitrogen  Tetroxide.  —  This  substance  is  liber- 
ated by  heating  nitrates,  other  than  those  of  potassium,  sodium,  or 
ammonium : 

2Cu(N08)2  -»  2CuO  +  4NO2  +  02. 


OXIDES  AND   OXYGEN  ACIDS  OF  NITROGEN  445 

In  most  cases  the  oxide  of  the  metal  remains.  When  the  mixed  gases 
are  led  through  a  U-tube  immersed  in  a  freezing  mixture,  the  tetroxide 
condenses  as  a  pale-yellow  liquid,  and  the  oxygen  passes  on. 

The  compound  may  also  be  made  by  direct  union  of  nitric  oxide 
and  oxygen,  or  by  oxidation  of  nitric  oxide  by  concentrated  nitric 
acid  (p.  443).  It  is  likewise  almost  the  sole  product  of  the  inter- 
action of  concentrated  nitric  acid  and  copper  (see  below).  If  any 
nitric  oxide  were  produced  by  the  primary  action,  it  would  be  oxidized 
to  nitrogen  tetroxide  in  passing  up  through  the  acid. 

Properties  of  Nitrogen  Tetroxide.  —  The  most  striking  pecul- 
iarity of  this  gas  is  that,  when  hot,  it  is  deep  brown  in  color,  and 
when  cold,  pale  yellow.  When  cooled,  it  gives  a  pale-yellow  liquid  boil- 
ing at  22°,  and  an  almost  colorless  solid  melting  at  about  —  12°.  The 
density  of  the  vapor  decreases  very  rapidly  from  27°  to  140°,  and 
increases  again  as  the  temperature  falls.  The  molecular  weights 
calculated  from  these  observations  are  :  at  27°,  76.7  ;  at  70°, 
55.6  ;  at  135°,  46.3  ;  at  154°,  45.7.  Now  the  molecular  weights  cor- 
responding to  the  formulae  N204  and  N02  are  92  and  46  respectively, 
so  that  these  results  mean  that  the  deep-brown  gas  is  N02,  and  that  as 
this  is  cooled  it  combines  to  form  the  colorless  N2O4.  Measurement  of 
the  depression  the  substance  causes  in  the  freezing-point  (cf.  p.  291)  of 
glacial  acetic  acid  gives  the  molecular  weight  92,  so  that  in  solution 
and  at  the  temperature  of  freezing  acetic  acid  (below  17°)  the  sub- 
stance is  all  N204. 

When  the  temperature  is  carried  above  154°,  by  passing  the 
brown  gas  through  a  red-hot  tube,  the  brown  color  disappears  once 
more,  and  nitric  oxide  and  oxygen  are  formed.  On  cooling,  the  same 
steps  through  brown  gas  to  pale-yellow  gas  are  retraced  : 


2NO  +  02  <=>  2N02  +±  N204, 
Colorless         Brown     Colorless. 

Since  nitrogen  tetroxide  yields  free  oxygen  more  readily  than  does 
nitric  oxide,  most  ordinary  combustibles  burn  in  it.  It  has  powerful 
oxidizing  properties  ;  and  "  fuming  nitric  acid,"  which  contains  it  in 
solution,  is  employed  when  oxidation  is  the  special  object  in  view. 

This  oxide  is  intermediate  in  composition  between  nitrous  and 
nitric  anhydrides,  and,  when  dissolved  in  cold  water,  gives  both  nitric 
and  nitrous  acids  : 

N,O4  +  H20  -*  HN08  +  HNO2. 


446  INORGANIC   CHEMISTRY 

If  a  base  is  present,  a  mixture  of  the  nitrate  and  nitrite  of  the  metal 
is  produced  (cf.  p.  275). 

When  the  water  is  not  cooled,  the  nitrous  acid  (y.v.),  being  un- 
stable, gives  nitric  oxide  and  nitric  acid,  so  that  the  result  is : 

3N02  +  H20  <=±  2HN08  +  NO. 

OXIDIZING  ACTIONS  OF  NITRIC  ACID. 

When  nitric  acid  gives  up  oxygen  to  any  body,  it  is  itself  reduced. 
Hence,  according  to  convenience,  we  shall  refer  to  oxidations  by,  or 
reductions  of  nitric  acid. 

Nascent  Hydrogen*  —  The  extent  to  which  the  acid  is  reduced 
by  nascent  hydrogen  depends  on  the  particular  metal  with  which  the 
hydrogen  is  in  contact  when  liberated  (cf.  p.  424).  Thus,  with  zinc 
and  very  dilute  nitric  acid,  almost  the  only  product,  aside  from  zinc 
nitrate,  is  ammonia : 

4Zn  +    8HN08  ->  4Zn(N08)2  (+  8H)  (1) 

(8H)  +      HN08  -»  NH8  +  3H20  (2) 

NH8  +      HN08  ->  NH4N08  (3) 
4Zn  +  10HN08  ->  4Zn  (N08)2  +  NH4NO8  +  311,0 

With  the  excess  of  nitric  acid,  ammonium  nitrate  is  formed.  Again, 
when  tin  is  the  metal,  the  reduction  is  less  complete,  and  hydroxy la- 
mine  is  a  product  (p.  423). 

When  metals  more  active  than  zinc,  such  as  magnesium,  are  used, 
some  of  the  hydrogen  escapes  oxidation  and  is  liberated. 

Heavy  Metals.  —  Metals  less  active  than  tin,  such  as  copper 
and  silver,  do  not  displace  hydrogen  from  dilute  acids  (p.  362),  but 
reduce  nitric  acid,  nevertheless,  and  are  converted  into  nitrates 
Platinum  and  gold  (cf.  p.  389)  alone  are  not  attacked.  Thus,  copper, 
with  somewhat  diluted  nitric  acid,  gives  cupric  nitrate  and  nitric  oxide 
(NO).  In  making  the  equation  for  this  action  we  may  resolve  the 
formula  of  nitric  acid  into  those  of  water  and  the  anhydride  H20,N2O6. 
This  shows  that  the  two  molecules  of  the  acid  will  give  2NO,  and  30 
will  remain : 

2HN08  ->  H20  +  2NO  ( +  3O)  (1) 

(3O)  +  6HN03  +  3Cu  -»  3H2O  +  3Gu(NOa),  (2) 

8HN08  +  3Cu  -» 4HjO  +  2NO  +  3Cu(N08)2 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  447 

The  nitric  oxide  is  liberated  as  a  colorless  gas,  but  forms  the  brown 
tetroxide  at  onoe  on  meeting  the  oxygen  of  the  air  (p.  443). 

When  concentrated  nitric  acid  is  used  with  copper,  almost  pure 
nitrogen  tetroxide  is  obtained  : 


2HNOS  -»  H20  +  2N02  (+0)  (1) 

(O)  +  2HN08  +  Cu  -»  H^Q  +  Cu  (N08)2  _  (2) 

4HN03  +  Cu  ->  211,0  +  2N02  +  Cu(N08)2 


The  equations  for  actions  like  the  above  may  be  built  up  from  partial  equations 
of  various  kinds  (c/.  p.  229).  Thus  we  may  begin  by  forming  the  nitrate  of  the 
metal,  and  then  use  the  balance,  consisting  of  hydrogen,  along  with  other  mole- 
cules of  nitric  acid  to  secure  the  oxide  and  water  : 

Cu  +  2HNO3  —  >  Cu(NO3)2(+  2H) 
(2H)  +  2HNO3-»  2H20  +  2NO2 


Cu  +  4HNO3  — »  Cu(N03)2+2H2O  + 2NO2 

As  the  subdivision  is  purely  arithmetical  (p.  229),  this  procedure  does  not  involve 
the  assumption  that  copper  does  actually  displace  hydrogen  as  a  free  element. 
Yet  it  would  not  necessarily  be  incorrect  to  make  even  this  supposition.  Although 
unable  to  liberate  hydrogen  in  quantity  from  a  dilute  acid,  copper  may  be  held, 
if  we  choose,  to  displace  a  minute  amount  of  it : 

Cu  +  2HN03  ±3  Cu(N03)2  +  H2,   or  Cu  +  2H'  ±5  Cu  "  +  H2, 

and  to  be  restrained  by  the  much  more  vigorous  reverse  action  (p.  362)  from  con- 
tinuing this  operation.  In  this  point  of  view  the  oxidation  of  the  trace  of  free 
hydrogen  by  the  excess  of  nitric  acid  continuously  annihilates  the  possibility  of 
reverse  action. 

Complexities  of  Oxidation  by  Nitric  Acid.  — The  above  are 
types  of  the  interactions  of  metals  with  nitric  acid.  In  actual  experi- 
ments the  behavior  is  usually  more  complex.  Thus,  as  a  rule,  the 
action  is  very  slow  at  first,  and  gathers  speed  with  the  accumulation  of 
the  reduction  products,  which  act  catalytically. 

Again,  different  concentrations  of  nitric  acid  give  different  prod- 
ucts with  the  same  metal.  The  most  conspicuous  effect  of  this  kind 
is  the  production  of  nitric  oxide  with  diluted  acid,  and  the  invariable 
formation  of  nitrogen  tetroxide  with  concentrated  acid.  This  is  ex- 
plained by  the  fact  that  nitrogen  tetroxide  cannot  pass  unchanged 
through  a  liquid  containing  much  water,  for  it  gives  nitric  acid  and 
nitric  oxide  with  the  latter  (p.  439).  Conversely,  where  the  nitric 
acid  is  concentrated,  nitric  oxide,  even  if  formed  by  the  interaction 
with  the  metal,  must  be  oxidized  to  nitrogen  tetroxide  as  it  passes  up 
through  the  liquid  (p.  443). 


448  INORGANIC   CHEMISTRY 

Finally,  intermediate  concentrations  give  mixtures  of  these  two 
oxides,  and,  with  zinc,  even  nitrous  oxide  (N20)  and  nitrogen  may  be 
found  in  considerable  quantities  in  the  gases  evolved. 

Non-  Metals.  —  With  non-metals  the  actions  are  different  in  so 
far  that  these  elements  do  not  give  nitrates.  Thus,  sulphur  boiled 
in  nitric  acid  gives  sulphuric  acid,  along  with  nitric  oxide,  equation  (3), 
or  with  nitrogen  tetroxide,  equation  (6),  or  with  both,  according  to  the 
concentration  of  the  acid  : 

2HN03  -4  2NO  +  H20  (+  30)  (1) 

(30)  +  HaO  +  S  -»  H2S04         _  (2) 

2HNO8  +  S  ->  2NO  +  H2S04  (3) 

2HN08  -*  2N02  +  H20  +  O      x  3  (4) 

(30)  +  H20  +  S  ->  H2S04  _  (5) 

S  ->  6NO2  +  2H2O  +  H2SO4  (6) 


The  reader  will  note  (cf.  p.  272)  that  a  separate  equation,  (3)  and  (6), 
must  be  made  for  the  formation  of  each  reduction  product.  If  NO 
and  N02  are  both  formed,  they  cannot  arise  from  the  same  molecule  of 
nitric  acid.  They  result  from  two  actions  which  are  independent, 
although  proceeding  simultaneously  in  the  same  vessel  (cf.  p.  231). 
Thus  the  equation  : 

2HN08  +  C  ->  H20  +  C02  +  NO  +  N02 


2, 


is  a  misrepresentation.  It  implies  that  equimolar  quantities  of  the 
two  oxides  of  nitrogen  are  formed.  But  this  could  only  occur  by 
chance,  and  the  balance  would  be  destroyed  the  next  moment  by  the 
lowering  in  the  concentration  of  the  acid,  giving  the  advantage  to  the 
nitric  oxide. 

Compounds,  —  Compounds  like  hydrogen  sulphide,  hydrogen 
iodide,  and  sulphurous  acid,  which  are  easily  oxidized,  interact  with 
nitric  acid.  With  diluted  nitric  acid  the  products  are  free  sulphur, 
iodine,  and  sulphuric  acid  respectively. 

The  mixture  of  nitric  acid  and  hydrochloric  acid  is  known  as 
aqua  regia.  The  chlorine  set  free  by  the  oxidation  of  the  hydro- 
chloric acid  is  more  active  than  is  the  ordinary  solution  of  chlorine  in 
water,  perhaps  in  consequence  of  catalytic  action  of  the  substances  in 
this  solution,  and  combines  with  gold  and  platinum  (q.v.),  converting 
them  into  chlorides.  Nitrosyl  chloride  (NOC1),  which,  however,  does 


OXIDES   AND   OXYGEN  ACIDS   OF  NITROGEN  449 

not  interact  directly  with   the  noble  metals,  is  formed  at  the  same 
time  : 

O 

C1-+2H0  +  01  +  01-^=0. 


The  interaction  with  platinum  is,  therefore  : 

8HC1  +  2HN03  +  Pt  ->  H2Pt016  +  2NOC1  +  4  H20. 

Chlorine,  in  this  active  state  of  the  element,  may  be  spoken  of  as 
nascent  chlorine  (cf.  p.  423). 

NITROUS  ACID,  HYPONITROUS  ACID,  AND  THEIR  ANHYDRIDES. 

Nitrites.  —  When  the  nitrates  of  potassium  and  sodium  are  heated, 
they  lose  one  unit  of  oxygen,  and  the  nitrites  remain  : 

2NaN03->  2NaN02  4-  O2. 

Commonly  lead  is  stirred  with  the  melted  nitrate  and  assists  in  the 
removal  of  the  oxygen.  The  litharge  (PbO)  which  is  formed  remains 
as  a  residue  when  the  sodium  nitrite  is  dissolved  for  recrystallization. 

Nitrous  Acid.  —  When  an  acid  is  added  to  a  dilute  solution  of 
a  nitrite,  a  pale-blue  solution  containing  nitrous  acid  is  obtained.  The 
acid  is  very  unstable,  however,  and,  when  the  solution  is  warmed, 
it  decomposes  : 

3HN02-»  HNO8  +  2NO  +  H2O. 

When  a  concentrated  solution  of  sodium  nitrite  is  acidified,  the 
nitrous  acid  decomposes  at  once,  and  a  brown  gas  containing  the 
anhydride  escapes  : 

2H*  +  2N02'  *=»  2HN02  <=»  H20  +  N2O3  \. 

This  behavior  distinguishes  a  nitrite  from  a  nitrate. 

Keducing  agents  deprive  nitrous  acid  of  part  or  all  of  its  oxygen  : 

2HI  +  2HN02  ->  2H20  +  2NO  +  I2. 

indigo  is  also  converted  by  it  into  isatin  (cf.  p.  269).  On  the  other 
hand,  oxidizing  agents  which  are  sufficiently  active,  like  acidified 
potassium  permanganate,  convert  nitrous  acid  into  nitric  acid  : 


450  INORGANIC  CHEMISTRY 

3H2S04  +  2KMn04  -4  K2S04  +  2MnS04  +  3H20  (  +  50)        (1) 
(50)  +  5Hfl  02  ->  5HN03  (2) 

3H2S04  +  2KMn04  +  5HN02  ->  KjSO4  -f  2MnSO4  +  3H2O  +  5HN03 


Nitrous  acid  is  much  used  in  the  making  of  organic  dyes. 

Nitrous  Anhydride.  —  A  study  of  the  gas  arising  from  the  decom- 
position of  nitrous  acid  shows  that  in  the  gaseous  state  the  anhydride 
is  almost  entirely  dissociated  : 


When  the  mixture  is  led  through  a  U-tube  immersed  in  a  freezing 
mixture  at  —  21°,  a  deep-blue  liquid  is  obtained  which  appears  to  be 
the  anhydride  itself.  This  begins  to  dissociate  before  reaching  its 
boiling-point,  and  at  +  2°  gives  off  nitric  oxide. 

The  same  equimolar  mixture  of  the  two  gases  is  obtained  by  the 
action  of  water  on  nitrosylsulphuric  acid  (p.  383). 

Hyponitrous  Acid.  —  This  acid  is  formed  by  the  interaction  of 
hydroxylamine  and  nitrous  acid  in  aqueous  solution  : 


H-0-N!H2-f-OiN-0-H->H20  +  H-O-N=N-0-H. 

With  nitrate  of  silver,  the  yellow,  insoluble  silver  hyponitrite  Ag2N202 
is  precipitated.  When  this  salt  is  shaken  with  an  ethereal  solution  of 
hydrogen  chloride,  the  acid  is  liberated,  and  the  insoluble  silver 
chloride  may  be  separated  by  filtration.  Finally,  evaporation  of  the 
ethereal  solution  leaves  hyponitrous  acid  as  a  white  mass.  It  explodes 
when  heated,  and  its  solution  in  water  is  an  exceedingly  feeble  acid. 
The  warm  aqueous  solution  decomposes  slowly,  giving  nitrous  oxide  : 


and  this  change  is  not  capable  of  reversal. 

Nitrous  O&ide.  —  Nitrous  oxide  is  prepared  by  heating  ammonium 
nitrate,  or  a  mixture  of  a  salt  of  ammonium  and  a  nitrate  : 

NH4N08  i*  2H20  +  N20. 

The  steam  condenses,  and  the  nitrous  oxide  may  be  collected  over 
warm  water,  or  dried  and  compressed  into  steel  cylinders. 

Its  solubility  in  cold  water  is  considerable  :  130  volumes  in  100  at 


OXIDES   AND   OXYGEN   ACIDS   OF  NITROGEN  451 

0°.  At  25°  this  falls  to  60  volumes  in  100.  In  dissolving,  the  gas 
forms  no  compound  with  water.  The  substance  melts  at  —  102.3°, 
and  boils  at  —  89.8°.  The  vapor  tension  of  the  liquid  at  0°  is  30.75 
atmospheres  ;  at  12°,  41.2  atmospheres  ;  and  at  20°,  49.4  atmospheres. 
The  critical  temperature  is  38.8°. 

A  glowing  splinter  of  wood  bursts  into  flame  when  introduced  into 
nitrous  oxide,  and  phosphorus,  sulphur,  and  other  combustibles,  burn 
in  it  with  much  the  same  vigor  as  in  oxygen.  In  all  cases  oxides  are 
formed,  and  nitrogen  is  set  free.  The  rapidity  with  which  bodies 
combine  with  oxygen  obtained  from  nitrous  oxide  is  doubtless  due  to 
the  fact  that  it  is  an  endothermal  compound,  and  the  heat  liberated  by 
its  decomposition  assists  the  ensuing  combustion  : 

2N20  ->  2N2  +  02  +  2  x  18,000  cal. 

It  is  to  be  noted  that  the  effect  of  the  heat  of  decomposition  will  be 
partly  offset  by  the  dilution  of  the  oxygen  with  nitrogen.  Yet  the 
proportion  of  nitrogen  to  oxygen  is  only  half  as  great  as  in  air,  so 
that  on  the  whole  the  conditions  are  much  more  favorable  to  combus- 
tion in  this  gas. 

Nitrous  oxide,  when  cold,  does  not  behave  like  free  oxygen.  Nitric 
oxide,  when  mixed  with  it,  gives  none  of  the  red  nitrogen  tetroxide. 
Metals  do  not  rust  in  it,  and  the  haemoglobin  of  the  blood  is  unable 
to  use  it  as  a  source  of  oxygen.  It  was  Davy  who  first  observed  that 
nitrous  oxide  could  be  taken  into  the  lungs,  and  that,  since  it  furnished 
no  oxygen,  insensibility  followed  its  use.  By  suitable  admixture  of  an 
amount  of  air  sufficient  to  sustain  life,  it  is  employed  as  an  anaesthetic 
for  minor  operations.  The  hysterical  symptoms  which  were  observed 
to  accompany  its  administration  caused  it  to  receive  the  name  of 
"laughing  gas." 

Graphic  Formulas  of  Nitric  Acid  and  its  Derivatives  :  Explo- 

sives. —  The  following  equation  shows  the  graphic  formulas  of  nitric 
acid  and  of  ammonium  nitrate  : 

H\  H\ 


N-OH  +  H-0-N-»N--0-N          +H20. 
^ 


0 


The  structural  formula  of  the  latter  is  intended  to  explain  the  fact 
that  the  salt  is  able  to  exist  at  all,  by  representing  the  oxygen  and 
hydrogen  as  being  separated  from  one  another  and  attached  to  differ- 


452  INORGANIC   CHEMISTRY 

ent  nitrogen  units.  When  the  equilibrium  of  the  system  is  disturbed 
by  heating,  the  oxygen  and  hydrogen  unite  to  form  water,  an  arrange- 
ment which  is  much  more  stable,  and  nitrous  oxide  (see  above)  escapes 
with  the  steam. 

The  decomposition  of  nitroglycerine  and  gun-cotton  (p.  441),  as 
well  as  of  ammonium  nitrite  (p.  416),  is  explained  in  the  same  way. 
These  substances  are  made  by  actions  which,  like  the  above  neutraliza- 
tion, take  place  in  the  cold,  and  the  groups,  containing  the  oxygen  on 
the  one  hand  and  carbon  and  hydrogen  on  the  other,  become  quietly 
united  without  more  serious  interaction.  Thus  the  formation  of  nitro- 
glycerine (p.  441)  appears  as  follows : 

H  H 

I  I 

H  -  C  -  OH  HO  -  N02  H  -  C  -  O  -  N02 

I  I 

H  -  C  -  OH  +  HO  -  N02  -»  H  -  C  -  0  -  N02  +  3H2O 

I  I 

H  -  C  -  OH      HO  -  NO2  H  -  C  -  O  -  N02 

I  I 

H  H 

When  the  nitroglycerine  is  heated,  or  receives  a  mechanical  shock, 
the  oxygen  all  unites  with  the  carbon  and  hydrogen,  and  the  nitrogen 
escapes : 

4C8H6(ON02)3-»  12C02  +  10H20  +  6N2  +  02. 

That  nitroglycerine  is  a  more  sensitive  explosive  than  gunpowder 
is  due  to  the  fact  that,  in  the  former,  the  materials  required  for  the 
chemical  change  are  already  within  the  same  molecule,  whereas  in 
the  latter  (q.v.)  they  are  contained  in  the  separate  molecules  of  a  mix- 
ture. Even  after  the  most  careful  incorporation,  the  oxygen  of  the 
potassium  nitrate  can  hardly  be  uniformly  so  near  to  the  carbon 
mechanically  mixed  with  the  salt,  as  are  these  elements  in  nitro- 
glycerine or  gun-cotton.  In  the  latter  the  oxidation  of  the  hydrogen 
and  carbon  is  intramolecular. 

Substances  like  hydrazoic  acid  (p.  422)  and  nitrogen  iodide  (p.  425) 
might  seem  to  constitute  a  third  kind  of  explosive.  Here  the  change 
consists  in  the  resolution  of  the  compound  into  its  constituents.  Still,  if 
we  consider  the  case  of  hydrazoic  acid,  for  example :  2K8H  — >  SN2  +  H2, 
we  see  that  the  action  consists,  after  all,  in  the  union  of  the  con- 
stituents to  form  the  more  stable  combinations  N2  and  H2.  It  is, 
therefore,  similar  in  principle  to  the  explosion  of  nitroglycerine. 


OXIDES   AND    OXYGEN   ACIDS   OF   NITROGEN  453 

The  Principle  of  Transformation  by  Steps.  —  It  may  have 
occurred  to  the  reader  as  strange  that  it  should  be  possible  to  make 
nitric  anhydride  by  distilling  a  warm  mixture  (p.  440)  when  the  product 
decomposes  spontaneously,  even  when  kept  in  the  cold.  How  can  a 
compound  be  fitted  together  under  certain  conditions,  when  under  the 
same  or,  even,  under  more  favorable  conditions  it  proves  to  be  incapable 
of  continued  existence  ?  We  should  expect  rather  that  obtaining  the 
products  of  its  decomposition  would  have  been  the  only  result  of  the 
effort  to  make  it. 

Extraordinary  as  this  fact  appears  to  be,  it  is  nevertheless  very 
commonly  encountered.  Perchloric  acid  is  made  by  a  distillation 
(p.  276)  and  afterwards  breaks  up  of  its  own  accord.  So,  also,  hypo- 
chlorites  are  formed  first  and  can  be  isolated.  But  under  the  same  con- 
ditions the  further  transformation  to  chlorates  will  occur  (p.  272).  A 
simple  case  is  that  of  sulphur  made  by  precipitation  (p.  376)  at  the  ordi- 
nary temperature.  Although  it  is  naturally  solid  below  119°,  yet, 
when  first  thrown  down,  it  is  in  the  form  of  liquid  droplets  which,  if 
undisturbed,  may  remain  fluid  for  weeks.  Similarly,  sulphur  vapor 
condenses  on  glass  in  drops  which  remain  liquid  until  they  are  touched 
or  rubbed.  Finally,  a  supersaturated  solution  (p.  159)  is  not  unlike 
cold  liquid  sulphur. 

In  all  these  cases  there  is  a  possibility  of  further  change,  which, 
when  it  comes,  will  liberate  heat  or  some  other  form  of  energy.  Thus, 
heat  is  set  free  when  the  liquid  sulphur  is  precipitated.  The  amount 
of  heat  would  have  been  greater,  by  the  heat  of  fusion  of  the  sulphur, 
if  solidification  had  occurred  simultaneously.  But,  in  spite  of  the 
existence  of  this  justification  for  the  final  step,  this  step  is  not  taken. 
So,  the  decomposition  of  the  vapor  of  the  perchloric  acid  or  of  the 
nitric  anhydride  would  have  added  to  the  amount  of  energy  liberated 
as  heat,  but  this  additional  step  was  postponed.  In  other  words, 
transformations  "which  proceed  spontaneously  and  with  evolution  of 
heat  may  go  forward  by  steps,  when  there  are  intermediate  substances 
capable  of  existence.  This  is  known  as  the  principle  of  transformation 
by  steps,  and  was  first  described  by  Ostwald. 

Exercises*  —  1.  Make  the  equation  for  the  interaction  of  ferrous 
chloride,  hydrochloric  acid,  and  nitric  acid  (p.  442),  and  for  all  the  ac- 
tions concerned  when  the  test  for  a  nitrate  (p.  443)  is  applied  to  sodium 
nitrate.  What  volume  (at  0°  and  760  mm.)  of  NO  is  obtained  from 
tme  formula-weight  of  nitric  acid  (p.  216,  Ex.  8)? 


454  INORGANIC   CHEMISTRY 

2.  Should  you  classify  as  molecular  compounds  (p.  443)  :    Chlorine 
hydrate,  ammonium  hydroxide,  KI8  (p.  235),  sulphurous  acid,  sodium 
pentasulphide  (p.  376)  ?     Justify  your  answer. 

3.  At  27°,  what  proportions  of  the  molecules  of  nitrogen  tetroxide 
are  in  the  forms  of  N02  and  N204  respectively  (p.  445)  ?  At  the  same 
temperature  what  fraction  of  the  material,  by  weight,  is  in  the  former 
condition  ?     What  are  the  relative  volumes  of  the  tetroxide,  and  of  the 
nitric  oxide  and  oxygen  obtained  by  its  decomposition  (p.  445)  ? 

4.  Make  an  equation  showing  the  production  of  nitrous  oxide  by  the 
action  of  zinc  on  nitric  acid. 

5.  Make  the  correct  equations  showing  the  formation  of  nitric  oxide 
and  nitrogen  tetroxide  by  the  interaction  of   carbon  and  nitric  acid 
(p.  448). 

6.  Justify  the  graphic  formula  assigned  to  nitric  acid  (p.  441). 

7.  Using  the  anhydride  method  (p.  274),  make  the  equations  for 
the  interactions  of  N2O4  and  water  (p.  445)  and  of  nitric  acid  and 
sulphur  (p.  448). 

8.  Make   a   classified   list,   with   examples,   of   all   the   kinds   of 
interactions  which,  in  this  and  preceding  chapters,  have  been  named 
oxidations  and  reductions  (e.g.  pp.  70,  72,  110,  170,  172,  231,  237,  269, 
306,  374,  420,  424,  446.     See  also  Chemistry  of  copper  and  tin). 


CHAPTER  XXVII 

PHOSPHORUS 

The  Chemical  Relations  of  the  Element.  —  There  are  many 
things  in  the  chemistry  of  phosphorus  and  its  compounds  which 
remind  us  of  nitrogen.  Yet  these  are  largely  referable  to  the  fact  that 
the  elements  are  both  non-metals  and  both  have  the  same  valences,  viz. 
three  and  five.  The  behavior  of  the  compounds  is  often  very  different. 
For  the  present  it  is  sufficient  to  say  that  both  give  compounds  with 
hydrogen,  NHS  and  PH3,  and  both  yield  oxides  of  the  forms  X203,  Xj04, 
and  X2O6.  The  first  and  last  of  these  oxides  are  acid-forming,  and 
phosphorus,  therefore,  gives  acids  corresponding  to  nitrous  and  nitric 
acids,  although  there  is  more  variety  in  the  proportion  of  water  com- 
bined with  the  anhydride  (cf.  p.  278).  The  element  is  thus  a  non-metal 
(see  Comparison  with  nitrogen  and  with  sulphur,  end  of  this  chapter). 

Occurrence.  —  This  element  is  found  widely  disseminated  in  na- 
ture, usually  in  the  form  of  phosphates.  Calcium  phosphate,  for  ex- 
ample, which  is  derived  (p.  278)  from  phosphoric  acid  (H8P04)  by  dis- 
placement of  its  hydrogen  by  calcium,  and  has,  therefore,  the  formula 
Ca^PO^a,  is  found  in  most  soils.  It  constitutes  a  large  part  of  the 
solid  material  of  the  bones  and  teeth  of  animals  and  of  the  beds  of 
fossil  bones  found  in  Florida  and  Tunis.  A  conspicuous  mineral 
related  to  this  substance  is  apatite,  Ca5F(P04)s.  It  is  found  in  large 
quantities  in  Canada,  and  is  a  component  of  many  rocks.  Complex 
organic  compounds  containing  phosphorus  are  essential  constituents  of 
protoplasm  and  of  the  materials  of  the  nerves  and  the  brain. 

Preparation.  —  Brand,  merchant  and  alchemist,  of  Hamburg, 
discovered  phosphorus  (1669)  by  distilling  the  residue  from  evapo- 
rated urine,  in  the  course  of  his  search  for  the  Philosopher's  stone. 
The  mode  of  preparing  it  from  bone-ash  was  first  published  by  Scheele 
(1771).  Green  bones  contain  about  58  per  cent  of  calcium  phosphate. 
After  the  gelatine  has  been  extracted  from  them,  by  means  of  water 
boiling  under  pressure,  they  are  subjected  to  destructive  distillation,  a 

465 


456  INORGANIC  CHEMISTRY 

process  which  yields  bone-oil.  The  residue  is  a  mixture  of  carbon 
(q.v.)  and  calcium  phosphate.  It  is  used  by  sugar  refiners  as  a  decol- 
orizer.  When  its  powers  in  this  direction  have  been  exhausted,  it  is 
calcined  —  that  is  to  say,  all  the  combustible  matter  is  burned  out  of  it, 
—  and  the  product  is  bone-ash.  Formerly  this  was  used  in  making 
phosphorus,  but  now  the  less  expensive  calcium  phosphate  of  fossil 
origin  is  employed. 

A  mixture  of  powdered  bone-ash  or  calcium  phosphate  and  sulphu- 
ric acid  (sp.  gr.  1.5  to  1.6)  is  heated  with  steam  and  stirred  in  a 
wooden  vat  : 

Ca3(PO4)2  +  3H2SO4  ->  2H8PO4  +  3CaS04. 

The  calcium  sulphate  is  partly  precipitated  during  the  heating.  The 
liquid  obtained  by  filtration  is  evaporated  in  leaden  pans.  During  this 
process  most  of  the  remainder  of  the  calcium  sulphate  is  deposited  and 
a  syrupy,  crude  phosphoric  acid  is  obtained.  This  acid  is  mixed  with 
sawdust,  or  carbon  in  some  form,  and  the  mixture  is  first  heated  to  a 
moderate  temperature  and  then  distilled  in  earthenware  retorts.  Two 
actions  take  place  in  succession.  The  phosphoric  acid  loses  water  and 
turns  into  metaphosphoric  acid,  then  the  latter  is  reduced  by  the  car- 
bon, carbon  monoxide  and  phosphorus  vapor  passing  off  : 


H8P04->H20 

2HP08  +  60  -+  H2  +  6CO  +  2P. 

A  white  heat  is  required  for  the  distillation,  and  a  pipe  from  the  tubu- 
lar clay  retort  conducts  the  vapors  into  cold  water,  in  which  the  phos- 
phorus collects. 

A  much  simpler  process  depends  on  the  use  of  the  electric  furnace 
(Fig.  90).  The  calcium  phosphate  is  mixed  with  the  proper  propor- 
tions of  carbon  and  silicon  dioxide  (sand),  and  the  mixture  is  intro- 
duced continuously  into  the  furnace.  The  discharge  of  an  alternating 
current  between  carbon  poles  produces  the  very  high  temperature 
which  the  action  requires.  The  calcium  silicate  which  is  formed  fuses 
to  a  slag,  and  can  be  withdrawn  at  intervals.  The  gaseous  products 
pass  off  through  a  pipe  and  the  phosphorus  is  caught  under  water  : 

Ca8(P04)2  +  3Si02  +  5C  ->  3CaSi08  +  5CO  -f-  2P. 

We  may  regard  the  phosphate  as  being  composed  of  two  oxides,  3CaO, 
P206.  It  thus  appears  that  the  calcium  oxide  has  united  with  the 


PHOSPHORUS 


457 


silica,  which  is  an  acid  anhydride  (cf.  p.  381)  :  CaO  +  Si02  — >  CaSiO3, 
while  the  phosphoric  anhydride  has  been  reduced. 

The  phosphorus,  after  purification,  is  cast  into  sticks  in  tubes  of 
tin  or  glass,  standing  in  cold  water. 

The  Electric  Furnace.  —  By  an  electric  furnace  is  understood  an 
electro-thermal  arrangement  in   which  the   heat  produced   by   some 
resistance   offered   to   the  current, 
such  as  that  of  an  air-gap  between 
the  carbons,  is  used  to  produce  chem- 
ical change.     Electrolysis  plays  no 
part   in    the    phenomena,   and    an1 
alternating  current,  which  can  pro-, 
duce  no  electrolytic  decomposition, 
is    generally    employed.       The   re- 
stricted area  within  which  the  heat 
is   developed    makes   possible    the 
attainment  of  a  high   temperature 
(see  Calcium  carbide). 

Physical  Properties.  —  There 
are  two  perfectly  distinct  kinds  of 
phosphorus,  known  as  ordinary,  or 
yellow  phosphorus,  and  red  phos- 
phorus. Yellow  phosphorus,  pre- 
pared as  described  above,  is  at  first 

transparent  and  colorless,  but  after  exposure  to  light  acquires  a 
superficial  coating  of  the  red  variety.  It  melts  at  44°  and  boils  at 
269°  (according  to  some  authorities,  at  287°).  Its  molecular  weight  is, 
at  313°,  128,  and  at  a  red  heat  119.8.  As  the  atomic  weight  is  31,  the 
formula,  within  this  range,  is  P4.  At  1700°  the  value  91.2  is  held  to 
indicate  partial  dissociation  into  P2.  In.  solution  the  formula  is  P4. 
Yellow  phosphorus  is  very  soluble  in.  carbon  disulphide,  less  soluble  in 
ether  and  other  organic  solvents,  and  insoluble  in  water.  It  is  ex- 
ceedingly poisonous,  less  than  0.15  g.  being  a  fatal  dose.  Continued 
exposure  to  its  vapor  causes  necrosis,  a  disease  from  which  match- 
makers are  liable  to  suffer.  The  jawbones  and  teeth  are  particularly 
liable  to  attack. 

Red  phosphorus  is  a  dull  red  powder  consisting  of  small  tabular 
crystals.  It  is  obtained  by  heating  yellow  phosphorus  to  about  250° 
in  a  vessel  from  which  air  is  excluded.  The  change  is  much  more 


FIG.  90. 


458  INORGANIC  CHEMISTRY 

rapid  at  slightly  higher  temperatures.  Since  a  great  amount  of  heat  is 
evolved  in  the  transformation,  the  action  is  apt  to  become  violent  and 
to  cause  volatilization  of  a  large  part  of  the  phosphorus.  In  presence 
of  a  trace  of  iodine  the  transformation  is  greatly  accelerated,  and  takes 
place  even  in  the  cold. 

Ked  phosphorus  does  not  melt,  but  passes  directly  into  vapor.  Its 
vapor  is  identical  with  that  of  yellow  phosphorus.  It  is  insoluble  in 
carbon  bisulphide  and  other  solvents.  It  is  not  poisonous,  and, 
unlike  yellow  phosphorus,  does  not  require  to  be  kept  under  water  to 
avoid  spontaneous  combustion. 

Chemical  Properties.  —  Yellow  phosphorus  unites  directly  with 
the  halogens  with  great  vigor.  It  unites  slowly  with  oxygen  in  the 
cold,  and  with  sulphur  and  many  metals  when  the  materials  are  heated 
together. 

The  slow  union  of  phosphorus  with  atmospheric  oxygen  is  accom- 
panied by  the  evolution  of  light,  although  the  temperature  is  not  such 
as  we  usually  associate  with  incandescence.  The  name  of  the  element 
(Gk.  <£<os,  light ;  <£epw,  to  bear)  records  this  property.  Apparently  the 
chemical  energy,  transformed  in  connection  with  the  oxidation,  is 
converted,  in  part  at  least,  into  radiant  energy  instead  of  completely 
into  heat.  *  A  curious  fact  in  connection  with  the  luminosity  and 
concomitant  oxidation  of  phosphorus  is  that  these  occurrences  depend 
upon  the  concentration  of  the  oxygen  gas  as  well  as  upon  the  tempera- 
ture. Thus,  phosphorus  does  not  shine  or  oxidize  in  pure  oxygen 
below  27°.  If  the  concentration  of  the  oxygen  is  reduced  to  200  mm. 
or  less  by  means  of  a  pump,  or  by  mixing  with  an  indifferent  gas  such 
as  nitrogen,  phosphorescence  becomes  perceptible  at  the  ordinary 
temperature.  This  explains  the  luminosity  shown  in  the  air.  At 
lower  temperatures,  lower  pressures  have  to  be  used.  The  phosphores- 
cence may  be  destroyed  by  the  vapor  of  turpentine  and  other  substances. 
All  these  phenomena  are  probably  due  to  the  intermediate  formation 
of  phosphorus  trioxide,  the  vapor  of  which  shows  the  same  effects. 

*  The  same  production  of  light  from  chemical  action  in  a  cold  body  is  seen  in 
the  luminosity  of  certain  parts  of  some  animals,  such  as  fireflies  and  some  species 
of  fish.  In  many  violent  chemical  changes  the  light  given  out  is  conspicuously 
greater  than  that  proper  to  the  temperature  produced  (cf.  p.  73),  and  must  come, 
therefore,  in  part,  directly  from  the  chemical  energy.  Thus,  burning  magnesium 
has  a  temperature  of  about  1350°,  while  the  production  of  light  of  the  same  char- 
acter, by  mere  incandescence,  would  require  a  temperature  of  about  6000°. 


PHOSPHORUS  459 

The  difference  in  behavior  of  pure  and  diluted  oxygen  may  be  shown  by  pour- 
ing absolution  of  phosphorus  in  carbon  disulphide  on  to  two  strips  of  filter  paper. 
One  of  the  strips,  hung  in  the  air,  catches  fire  as  soon  as  the  evaporation  of  the 
solvent  has  exposed  a  large  area  of  finely  divided  phosphorus.  The  other,  hung 
in  a  jar  of  oxygen,  remains  unaffected,  but  becomes  ignited  instantly  upon  removal 
from  the  jar. 

The  slow  oxidation  of  phosphorus  is  accompanied  by  the  produc- 
tion of  >  ozone,  but  the  nature  of  the  action  is  still  unknown. 

Chemical  Properties  of  Red  Phosphorus. — This  variety  of  the 
element,  since  it  is  formed  with  evolution  of  heat,  contains  less  energy 
than  yellow  phosphorus  and  is  much  less,  active.  It  does  not  catch 
fire  in  the  air  below  240°,  while  ordinary  phosphorus  ignites  at  35-45°. 
Indeed,  it  is  the  vapor  that  begins  to  combine  with  oxygen,  and 
this  behavior  is  only  an  independent  proof  of  the  low  vapor  tension  of 
the  red  variety.  When  the  vapor  tension  of  yellow  phosphorus  has 
reached  760  mm.  (at  269°,  the  b.-p.),  that  of  red  phosphorus  is 
almost  imperceptible.  Even  when  the  former  has  become  two  atmos- 
pheres, the  latter  is  still  very  small. 

Red  phosphorus  is  to  'be  regarded  as  the  normal,  stable  form  of 
phosphorus.  The  fact  that  yellow  phosphorus  can  be  kept  a  great 
length  of  time,  and  is  changed  but  slightly  on  exposure  to  light,  only 
shows  that  the  transformation  into  a  stabler  condition  is  retarded  by 
the  lowness  of  the  temperature  (cf.  p.  72).  The  relation  between  the 
two  varieties  of  phosphorus  is  quite  distinct  from  that  between  rhom- 
bic and  monoclinic  sulphur  (p.  368).  In  the  latter  case  there  is  a 
definite  temperature  of  transformation  (96°)  above  which  one  form 
completely  disappears,  and  below  which  the  other  form  is  incapable  of 
permanent  existence.  With  the  varieties  of  phosphorus  no  such  point 
of  transformation  exists.  The  red  phosphorus  is  the  more  stable  at 
all  temperatures  at  which  both  forms  are  known.  The  yellow  turns 
into  red,  but  the  red  never  into  the  yellow.  It  is  only  by  condensing 
the  vapor  that  the  yellow  kind  is  obtained.  It  seems  probable  that 
red  phosphorus  is  a  polymer  (p.  242)  of  yellow  phosphorus  (P4). 
The  production  from  the  vapor  (also  P4)  of  the  yellow  solid,  instead 
of  the  red  solid  whose  formation  would  be  accompanied  by  a  larger 
liberation  of  heat,  is  simply  an  illustration  of  the  principle  of  transfor- 
mation by  steps  (p.  453). 

When  two  or  more  forms  of  an  element,  or  even  of  a  compound  (see  Isomers), 
occur,  they  are  commonly  spoken  of  as  allotropic  modifications.  The  term  is 
applied  to  oxygen  and  ozone  (q.v.)  which  are  certainly,  and  to  red  and  yellow 


460  INORGANIC   CHEMISTRY 

phosphorus  which  are,  probably,  chemically  distinct  substances.  It  is  used  also  of 
rhombic  and  monoclinic  sulphur,  where  the  difference  is  purely  physical.  In  short, 
It  has  at  present  no  scientific  value,  for  it  covers  a  heterogeneous  mass  of  phe- 
nomena which,  in  part,  still  await  elucidation.  If  allotropic  modifications 
were  to  be  defined  as  substances  (p.  32)  composed  of  the  same  materials, 
but  possessing  different  proportions  of  free  or  available  energy,  and, 
therefore,  different  physical  properties  and  different  degrees  of  chemical  activity, 
which  is  apparently  the  sense  in  which  the  expression  is  commonly  employed,  then 
ice,  water,  and  steam  would  be  examples  of  such  substances.  In  each  of  the 
above  four  illustrations,  the  second  (in  the  case  of  water,  the  third)  is  the  more 
active  form. 

Uses  of  Phospliorus.  —  The  greater  part  of  the  phosphorus  of 
commerce  is  employed  in  the  manufacture  of  matches.  The  first  arti- 
cles of  this  sort  (1812)  were  sticks  coated  with  sulphur  and  tipped 
with  a  mixture  of  potassium  chlorate  and  sugar.  For  ignition  they 
were  dipped  into  a  bottle  containing  asbestos  moistened  with  concen- 
trated sulphuric  acid.  Matches  involving  the  use  of  phosphorus  (1827) 
have  now  displaced  all  others.  In  making  common  matches  which 
strike  on  any  rough  surface,  the  sticks  are  first  dipped  in  melted 
sulphur  or  paraffin  to  the  extent  of  about  half  an  inch.  The  head  is 
aften  composed  of  manganese  dioxide  or  red  lead,  and  a  little  potas- 
sium chlorate,  which  supply  oxygen,  a  small  proportion  of  free  phos- 
phorus and  antimony  trisulphide,  which  are  both  combustible,  and 
dextrine  or  glue.  A  paste  made  of  these  materials  is  spread  evenly 
upon  a  slab,  and  the  prepared  sticks  fixed  in  a  frame  are  dipped  once 
or  twice  in  the  mixture. 

In  the  case  of  "  safety  "  matches,  the  mixture  upon  the  head  is  not 
easily  ignited  by  itself.  It  is  composed  of  potassium  chlorate  or  di- 
chromate,  some  sulphur  or  antimony  trisulphide,  and  a  little  powdered 
glass  to  increase  the  friction,  all  held  together  with  glue.  Upon  the 
rubbing  surface  on  the  box  is  a  thin  layer  of  antimony  trisulphide 
mixed  with  red  phosphorus  and  glue.  The  friction  converts  a  little 
of  the  red  phosphorus  into  vapor.  To  prevent  smoldering  of  the 
burned  matches,  the  upper  ends  of  the  sticks  are  sometimes  soaked 
in  a  solution  of  alum  or  sodium  phosphate. 

A  small  amount  of  yellow  phosphorus  is  employed  in  making  rat 
poison,  which  is  a  mixture  of  phosphorus,  lard  (as  solvent),  and  flour, 
made  into  dough. 

Phosphine.  —  Three  hydrides  of  phosphorus  are  known.  These 
are,  phosphine  PH3  (a  gas),  a  liquid  hydride  P2H4,  which  is  presumably 


PHOSPHORUS  461 

the  analogue  of  hydrazine  (NaH4),  and  a  solid  hydride  P4H2.  Fhosphine 
does  not  seem  to  be  produced  under  ordinary  circumstances  by  the 
direct  union  of  the  elements.  It  is  formed  slowly,  however,  by  the 
action  of  nascent  hydrogen,  from  zinc  and  hydrochloric  acid  at  70°,  upon 
yellow  phosphorus.  The  gas  may  be  made  by  boiling  yellow  phos- 
phorus with  potassium  hydroxide  solution  in  an  apparatus  similar  to 
that  used  for  generating  hydrogen.  Potassium  hypophosphite  is 
formed  at  the  same  time  : 

3KOH  +  4P  +  3H20  ->  3KH2P02  +  PH3. 

The  gas  made  in  this  fashion  contains  a  small  proportion  of  the  vapor 
of  the  liquid  hydride,  which  is  spontaneously  inflammable,  and  conse- 
quently the  mixture  catches  fire  on  emerging  into  the  air  from  the 
delivery  tube.  To  avoid  explosions,  the  air  in  the  flask  must  be  dis- 
placed by  hydrogen  or  illuminating-gas  before  heat  is  applied.  This 
product  contains  also  -free  hydrogen,  in  increasing  quantities  as  the 
action  goes  on,  in  consequence  of  the  reduction  of  the  water  and  potas- 
sium hydroxide  by  the  potassium  hypophosphite  :  KH2PO2  -f-  KOH  -f- 
H2O  — >  K2HP04  +  2H2.  Potassium  phosphate  is  formed. 

The  simplest  method  of  preparing  the  gas  is  by  the  action  of  water 
upon  calcium  phosphide : 

Ca,P,.+  6H20  ->  3Ca(OH)2  +  2PH3. 

This  action  is  analogous  to  that  of  water  upon  magnesium  nitride 
(p.  417)  by  which  ammonia  is  produced.  In  consequence  of  the  fact 
that  calcium  phosphide  is  a  substance  of  irregular  composition,  a  mix- 
ture of  all  three  hydrides  is  generally  obtained.  By  passing  the  gas 
through  a  strongly  cooled  delivery  tube,  however,  the  liquid  compound 
is  condensed  and  fairly  pure  phosphine  passes  on. 

Phosphine  is  a  colorless  gas,  which  is  easily  decomposed  by  heat 
into  its  elements.  When  burned,  it  forms  phosphoric  acid.  It  is  exceed- 
ingly poisonous  and,  unlike  ammonia,  it  is  insoluble  in  water,  and  pro- 
duces no  basic  compound  corresponding  to  ammonium  hydroxide  when 
brought  in  contact  with  this  substance.  It  resembles  ammonia,  for- 
mally at  least,  in  uniting  with  the  hydrogen  halides  (see  below).  It 
differs  from  ammonia,  however,  inasmuch  as  it  does  not  unite  with  the 
oxygen  acids.  Phosphine  acts  upon  solutions  of  some  salts,  precipitat- 
ing phosphides  of  the  metals : 

3CuSO<  +  2PH8  ->  Cu8P3  +  3H2S04. 


462  INORGANIC   CHEMISTRY 

The  liquid  hydrogen  phosphide  boils  at  57°.  The  molecular  weight, 
as  determined  by  the  density  of  its  vapor,  shows  the  formula  to  be  P2H4. 
It  forms  no  salts,  and  is  therefore  quite  unlike  hydrazine.  When 
exposed  to  light,  it  decomposes,  giving  phosphine  and  the  solid  hydride. 

Phosphonium  Compounds.  —  Hydrogen  iodide  unites  with  phos- 
phine to  form  a  colorless  solid  crystallizing  in  beautiful,  highly  refract- 
ing, square  prisms  :  PH8  -f-  HI  —>  PH4I.  Hydrogen  chloride  combines 
similarly  with  phosphine,  but  only  when  the  gases  are  cooled  by  a  freez- 
ing mixture,  or  are  brought  together  under  a  total  pressure  of  18  atmos- 
pheres at  14°.  When  the  pressure  is  released,  rapid  dissociation  occurs. 
This  dissociation  is  one  of  the  many  cases  where  an  action  which 
absorbs  heat,  nevertheless  goes  on  spontaneously  (cf.  p.  27).  The 
indispensable  fall  in  the  energy  of  the  system  takes  place  by  virtue  of 
the  diffusion  of  the  constituents,  and  in  amount  this  more  than  offsets 
the  heat  acquired. 

In  imitation  of  the  ammonia  nomenclature,  these  substances  are 
called  phosphonium  iodide  and  phosphonium  chloride.  They  are 
entirely  different,  however,  from  the  corresponding  ammonium  deriva- 
tives, for  the  PH4'  ion  is  unstable.  When  brought  in  contact  with 
water  they  decompose  into  their  constituents,  the  hydrogen  halide 
going  into  solution,  and  the  phosphine  being  liberated  as  a  gas. 

HaUdes  of  Phosphorus.  —  The  existence  of  the  following  halides 
has  been  proved  conclusively : 

P2I4  (solid) 

PF8  (gas)         PC18  (liquid)         PBr8  (liquid)         PI3    (solid) 
PF6  (gas)         PC16  (solid)          PBr6  (solid)  

These  substances  may  all  be  formed  by  direct  union  of  the  elements. 
They  are  incomparably  more  stable  than  are  the  similar  compounds  of 
nitrogen.  They  are  all  decomposed  by  contact  with  water,  and  give  an 
oxygen  acid  of  phosphorus  and  the  hydrogen  halide  (see  below).  This 
action  was  used  in  the  preparation  of  hydrogen  bromide  (p.  231)  and 
hydrogen  iodide  (p.  238). 

Phosphorus  trichloride  is  made  by  passing  chlorine  gas  over 
melted  phosphorus  in  a  flask  until  the  proper  gain  in  weight  has 
occurred.  It  is  a  liquid,  boiling  at  76°.  When  excess  of  chlorine  is 
employed,  phosphorus  pentachloride,  which  is  a  white  solid  body,  is 


PHOSPHORUS  463 

formed.  When  moist  air  is  blown  over  any  of  these  substances,  the 
water  is  condensed  to  a  fog  by  the  hydrogen  halide.  In  the  case  of  the 
interaction  of  phosphorus  pentachloride  and  water,  phosphoric  acid  is 
formed  : 

PC15  +  4H20  _»  H3P04  +  5HC1. 

With  a  limited  supply  of  water  the  hydrolysis  is  not  so  complete, 
and  phosphoryl  chloride  *  (phosphorus  oxychloride),  a  liquid  boiling  at 
107°,  is  produced  ; 

PC15  -f  HOH  ->  POC13  +  2HC1. 

This  interaction  of  phosphorus  pentachloride  and  water  is  a  per- 
fectly general  one,  and  takes  place  with  most  compounds  containing 
hydroxyl.  Thus,  when  alcohol  (which  differs  from  water  in  having 
ethyl  (C2H5),  instead  of  hydrogen,  combined  with  hydroxyl)  is  poured 
upon  it,  ethyl  chloride,  phosphoryl  chloride,  and  hydrogen  chloride  are 
formed  : 

C2H5OH  +  PC15  ->  C2H5C1  +  POC13  +  HC1. 


The  same  action  takes  place  with  all  carbon  compounds  containing  hy- 
droxyl, and  is  used  as  a  means  of  showing  the  presence  of  this  group 
in  their  structure.  The  reaction  is  shown  by  inorganic  compounds 
also.  Thus,  anhydrous  sulphuric  acid  gives  sulphuryl  chloride,  which 
may  be  separated  from  the  phosphoryl  chloride  by  fractional  distilla- 
tion (see  Petroleum)  : 

S02(OH)2  +  2PC15  ->  80,012  +  2POCI,  +  2HC1. 

j 

Phosphorus  pentachloride,  when  heated,  reaches  a  vapor  tension 
of  760  mm.  at  140°,  and  while  still  solid.  It  therefore  passes  freely 
into  vapor  (boils,  so  to  speak)  at  this  temperature,  and  condenses  di- 
rectly to  the  solid  form.  This  sort  of  distillation  is  called  sublimation. 
At  a  pressure  above  that  of  the  atmosphere  it  melts  at  148°.  This  is 
simply  a  case  in  which  the  vapor  tension  of  the  solid,  increasing  with 
rise  in  temperature,  happens  to  pass  the  arbitrary  value  of  the  opposing 
pressure  (one  atmosphere)  peculiar  to  experiments  carried  on  in  open 
vessels,  before  the  melting-point  is  reached.  The  same  phenomenon  is 
shown  by  sulphur  trioxide  (p.  381). 

*  This  substance  is  a  mixed  anhydride  (p.  397)  of  phosphoric  acid  and  hydro- 
gen chloride. 


464  INORGANIC   CHEMISTRY 

Phosphorus    pentachloride  (cf.  p.  255)  and    pentabromide,  when 
vaporized,  are  partially  dissociated : 

PBr5  <±  PBr8  +  Bra. 

Since  the  first  two  members  of  this  equilibrium  are  colorless,  while  the 
bromine  is  brown,  this  action  may  be  used  to  illustrate  the  effect  upon 
a  system,  of  increasing  the  concentration  of  one  of  the  interacting 
substances  (p.  249).  Two  tubes,  of  equal  volume  and  containing 
equal  amounts  of  the  pentabromide  are  prepared.  'A  small  amount  of 
the  tribromide  is  added  to  the  second,  and  both  are  sealed  up.  When 
the  tubes  are  now  heated  to  the  same  temperature,  the  contents  of  the 
second  will  be  less  strongly  colored  by  bromine  in  consequence  of  the 
greater  activity  in  it  of  the  reversing  action. 


Oxides  of  Phosphorus*  —  The  oxides  of  phosphorus  are  the  so- 
called  trioxide  P4O6,  the  pentoxide  P206,  and  a  tetroxide  P204. 

The  pentoxide  is  a  white  powder  formed  when  phosphorus  is  burned 
with  a  free  supply  of  oxygen.  It  unites  with  water  with  great  violence 
to  form  metaphosphoric  acid  (see  below),  and  hence  is  known  as 
phosphoric  anhydride  :  P205  -f  H20  — »  2HP08.  In  the  laboratory 
this  action  is  frequently  utilized  for  drying  gases  (p.  101)  and  for  re- 
moving water  from  combination  (p.  276).  The  vapor  density  of 
the  pentoxide  indicates  that  its  formula  is  P4010,  but  its  whole  chemical 
behavior  is  equally  well  represented  by  the  simpler  formula. 

The  trioxide  is  obtained  by  burning  phosphorus  in  a  tube  with  a 
restricted  supply  of  air.  It  is  a  white  solid,  melting  at  22.5°  and  boil- 
ing at  173°.  On  account  of  the  ease  with  which  it  may  be  volatilized, 
it  can  be  separated  by  distillation  from  any  pentoxide  formed  at  the 
same  time.  The  operation  must  be  carried  out  in  an  apparatus  from 
which  the  air  is  excluded,  as  the  trioxide  unites  spontaneously  with 
oxygen.  The  vapor  density  of  the  substance  shows  that  its  formula  is 
P406.  This  formula  is  preferred  to  the  simpler  one  because,  although 
the  oxide  is  the  anhydride  of  phosphorous  acid,  it  nevertheless  unites 
exceedingly  slowly  with  cold  water  to  form  this  substance.  It  inter- 
acts vigorously  with  hot  water,  but  phosphine,  red  phosphorus,  hypo- 
phosphoric  acid,  and  phosphoric  acid  are  amongst  the  products,  and 
very  little  phosphorous  acid  escapes  decomposition.  When  this  oxide 
is  heated  to  440°  it  decomposes,  giving  the  compound  P204  and  red 
phosphorus. 


PHOSPHORUS  465 

Acids  of  Phosphorus.  —  There  are  four  different  acids  of 
phosphorus  in  which,  probably,  four  distinct  stages  of  oxidation  are 
shown.  The  highest  stage  is  represented  by  three  phosphoric  acids, 
where  the  degree  of  hydration  of  the  anhydride  varies.  These  are 
orthophosphoric  acid  H8P04,  pyrophosphoric  acid  H4P207,  and  metaphos- 
phoric  acid  HP08.  The  other  acids  are  less  important.  They  are 
phosphorous  acid  (H3P03),  hypophosphoric  'acid  (H4P206),  and  hypo- 
phosphorous  acid  (H8P02). 

The  Phosphoric  Acids.  —  The  relation  between  the  three  different 
phosphoric  acids  may  be  seen  by  considering  them  as  being  formed 
from  phosphorus  pentoxide  and  water.  It  will  be  remembered  that  in 
the  majority  of  cases  already  considered,  this  sort  of  action  takes  place 
for  the  most  part  in  but  one  way.  Thus,  nitric  acid  is  known  in  but  one 
form,  which  is  produced  by  the  union  of  one  molecule  each  of  nitrogen 
pentoxide  and  water  :  N206  +  H2O  ->  2HN03.  Similarly  the  chief  sul- 
phuric acid  is  the  one  formed  from  one  molecule  of  sulphur  trioxide 
and  one  molecule  of  water  :  SO8  +  H20  — » H2S04,  although  here  we 
have  both  the  hydrate  H2S04,  H20,  which  might  be  written  H4S06  and 
disulphuric  acid  H2S2O7.  Referred  to  the  anhydride  these  three  acids 
are  H20,S03,  2H20,S03,  and  H20,2S03.  Periodic  acid  (p.  278)  has 
a  set  of  even  more  complexly  related  acids  or  salts. 

Now,  when  phosphoric  anhydride  acts  upon  water  we  obtain  a  solu- 
tion which,  on  immediate  evaporation,  leaves  a  glassy  solid,  HP08, 
known  as  metaphosphoric  acid.  This  is  H20,  P206.  When,  however, 
the  solution  is  allowed  to  stand  for  some  days,  or  is  boiled  with  a  little 
dilute  nitric  acid  whose  hydrion  acts  catalytically,  the  residue  from 
evaporation  is  H8P04,  orthophosphoric  acid  : 

P206  +  3H20  ->  2H3P04     or     HP03  +  H20  ->  H8PO4. 

This  acid  is  3H2O,  PS06,  and  no  further  addition  of  water  can  be 
effected. 

Conversely,  when  orthophosphoric  acid  is  kept  at  about  255°  for  a 
time,  it  slowly  loses  water,  and  H4P207,  pyrophosphoric  acid,  is  ob- 
tained : 

2H8P04  ->  H2O  +  H4P2O7. 

This  acid  is  2H20,  P206.  Further  desiccation  leaves  metaphosphoric 
acid,  which  cannot  be  further  resolved  into  phosphorus  pentoxide  and 
water.  When  dissolved  in  water,  pyrophosphoric  acid  slowly  resumes 


466  INORGANIC  CHEMISTRY 

the  water  which  it  has  lost  and  gives  the  ortho-acid  again.  The  pyro- 
acid  does  not  seem  to  be  formed  by  hydration  of  the  meta-acid,  but 
only  by  dehydration  of  the  ortho-acid. 

The  relations  of  all  these  substances  are  more  clearly  seen  in  the 
graphic  formulas : 

=  0 

-o-H         r=o  r=o 

-0-H        P  •{  -O-H          P  -(  -0-H 
-0-H 


j  -O-H          P  4  -0-H  f  =0 

I  -0-H  [=0  P  <j  =0 

[     O  ->  <-          [     0 

r  =< 

{-: 


_0_H  f_0-H                r=0                 PJ  = 

-0-H  P  J  -0-H  p  J  -0-H 

-O-H  |  =0                        I  =0 
=0 

The  addition  or  removal  of  water  leaves  the  valence  of  the  phosphorus 
unchanged. 

Pyrosulphuric  acid  and  its  salts  when  dissolved  in  water  give  sul- 
phuric acid  and  acid  sulphates  respectively.  That  is  to  say,  the  ion  S207 
is  not  capable  of  existence.  But  the  very  slow  rate  at  which  the  less 
hydrated  phosphoric  acids  change  into  the  more  hydrated  ones  shows 
that  ions  like  P04'",  P03',  and  P2O7""  may  be  comparatively  stable. 
The  behavior  of  solutions  of  the  salts  shows  this  even  more  clearly. 

Orthophosphoric  Acid  and  Its  Salts.  —  As  we  have  seen,  ordi- 
nary calcium  phosphate  is  the  source  of  the  impure,  commercial  acid. 
Pure  orthophosphoric  acid  may  be  made  by  boiling  red  phosphorus 
with  slightly  diluted  nitric  acid  and  evaporating  the  water  and 
excess  of  nitric  acid.  The  product  is  a  white,  crystalline,  deliques- 
cent solid. 

This  acid  is  much  weaker  than  sulphuric  acid,  and  is  dissociated 
chiefly  into  the  ions  H*  and  H2PO/.  The  dihydrophosphanion  is 
broken  up  to  some  extent  into  H"  and  HP04",  as  we  learn  from  the 
fact  that  the  solution  of  the  sodium  salt  NaH2P04  is  acid.  The  ion 
HP04"  is  hardly  dissociated  at  all,  for  a  solution  of  the  salt  Na,jHP04 
is  alkaline  in  reaction. 

As  a  tribasic  acid,  it  forms  salts  of  three  kinds,  such  as  NaH2P04, 
Na2HP04,  and  Na8P04.  These  are  known  respectively  as  primary, 
secondary,  and  tertiary  sodium  orthophosphate.  The  primary  sodium 
phosphate  is  faintly  acid  in  reaction.  The  secondary  one  is  slightly 
alkaline,  because  of  hydrolysis  arising  from  the  tendency  of  the 


PHOSPHORUS  467 

hydrion  of  the  water  to  combine  with  the  HPO4"  to  form  H2PO/  (cf. 
p.  344).  The  tertiary  phosphate  is  stable  only  in  solid  form,  and  can 
be  made  by  evaporating  to  dryness  a  mixture  of  the  secondary 
phosphate  and  sodium  hydroxide  : 


Na2HP04  +  NaOH  <=»  N^PO,  +  H2O  f. 

When  the  product  is  dissolved  in  water,  the  action  is  reversed  (cf. 
p.  375).  Mixed  phosphates  are  also  known,  particularly  sodium-am- 
monium phosphate  (microcosmic  salt),  Na]STH4HP04,  and  the  insoluble 
magnesium-ammonium  phosphate,  MgNH4PO4. 

Primary  calcium  phosphate,  known  in  commerce  as  "  superphos- 
phate," is  used  as  a  fertilizer.  On  account  of  its  insolubility,  and 
since  plants  can  take  up  soluble  substances  only,  calcium  phosphate  is 
of  relatively  little  service  to  plants.  It  is  therefore  converted  into  the 
"  superphosphate,"  which  is  soluble,  by  treatment  with  dilute  sulphuric 
acid: 

Ca3(P04)2  +  2H2S04^±  2CaS04  +  CaH4(P04)2. 


The  tertiary  phosphates  are  unchanged  by  heating.  The  primary 
and  secondary  phosphates,  however,  retaining,  as  they  do,  some  of  the 
original  hydrogen  of  the  phosphoric  acid,  are  capable  of  losing  water 
like  phosphoric  acid  itself,  when  heated.  The  actions  are  slowly  re- 
versed when  the  products  are  dissolved  in  water : 

NaH2P04  ?=>  NaPO3  +  H2O, 
±  Na4P2O7  +  H2O. 


It  will  be  seen  that  the  meta-  and  pyrophosphates  of  sodium  are 
formed  by  these  actions  ;  and  this  is  indeed  the  simplest  way  of  form- 
ing these  substances,  since  the  acids  themselves  are  not  permanent  in 
solution,  and  are  too  feeble  to  lend  themselves  to  exact  neutralization. 
Ammonium  salts  of  phosphoric  acid  lose  ammonia,  as  well  as  water, 
when  heated  (cf.  p.  421).  Thus,  microcosmic  salt  gives  primary 
sodium  phosphate  : 

NaNH4HPO4->NaH2P04  +  NH8, 
and  this  in  turn  is  converted  into  the  metaphosphate  by  loss  of  water. 

Pyrophosphoric  Acid.  —  This  acid,  obtained  by  heating  ortho- 
phosphoric  acid,  may  be  prepared  in  pure  form  by  making  the  spar- 
ingly soluble  lead  salt  from  sodium  pyrophosphate,  and  precipitating 


468  INORGANIC   CHEMISTRY 

the  plumbion  by  addition  of  sulphuric  acid.  In  solution  it  gradually 
reunites  with  water.  Although  tetrabasic,  having  four  hydrogen 
atoms  which  may  be  displaced  by  metals,  only  two  kinds  of  salts 
of  this  acid  are  known.  These  are  the  normal  salts,  such  as  Na4P207, 
and  those  in  which  one-half  of  the  hydrogen  has  been  displaced  by 
a  metal  such  as  Na2H2P207. 

Metaphosphoric  Acid.  —  This  is  the  "  glacial  phosphoric  acid  " 
of  commerce,  and  is  usually  sold  in  the  form  of  transparent  sticks. 
It  is  obtained  by  heating  orthophosphoric  acid,  or  by  direct  union  of 
phosphorus  pentoxide  with  a  small  amount  of  cold  water.  It  passes 
into  vapor  at  a  high  temperature,  and  its  vapor  density  corresponds  to 
the  formula  (HP08)2.  The  existence  of  certain  complex  salts  confirms 
our  belief  in  the  existence  of  a  tendency  to  association  (p.  242). 

Sodium  metaphosphate  NaP08,  in  the  form  of  a  small  globule 
obtained  by  heating  microcosmic  salt  on  a  platinum  wire,  is  used  in 
analysis.  When  minute  traces  of  oxides  of  certain  metals  are 
placed  upon  such  a  globule,  known  as  a  bead,  and  heated  in  the  Bun- 
sen  flame,  the  mass  is  colored  in  various  tints  according  to  the  oxide 
used  (bead  teat).  This  action  may  be  understood  when  we  consider 
that  sodium  metaphosphate  takes  up  water  to  form  primary  sodium 
orthophosphate  :  NaPO8  -h  H20  — >  NaH2P04.  In  the  same  way,  but  at 
higher  temperatures,  it  is  able  to  take  up  oxides  of  elements  other  than 
hydrogen,  giving  mixed  orthophosphates.  Thus  with  oxide  of  copper 
a  part  of  the  metaphosphate  unites  according  to  the  equation  : 

NaP08  +  CuO  ->NaCuP04, 
and  the  product  confers  a  green  tinge  on  the  bead. 

Distinguishing  Tests.  —  When  a  solution  of  nitrate  of  silver  is 
added  to  a  solution  of  orthophosphoric  acid  or  any  soluble  orthophos- 
phate, a  yellow  precipitate  of  silver  orthophosphate  Ag3P04  is  produced. 
This  is  a  test  for  phosphanion.  With  pyrophosphoric  acid  or  any  pyro- 
phosphate  the  product  is  white  Ag4P2O7.  With  metaphosphoric  acid  a 
white  precipitate,  AgP08,  is  obtained  also.  Metaphosphoric  acid  coagu- 
lates a  clear  solution  of  albumen,  while  pyrophosphoric  acid  has  no 
visible  effect  upon  it. 

A  test  for  orthophosphoric  acid,  or  rather  the  ion  PO/",  consists  in  adding  a 
drop  of  the  solution  containing  this  ion  to  a  solution  of  ammonium  inolybdate  (q.v.) 
in  dilute  nitric  acid.  A  copious  yellow  precipitate  of  an  ammonium  phospho- 


PHOSPHORUS  469 

molybdate,  (NH4) ,PO4,llMoO3,  6H2O,  appears  on  warming.  In  presence  of  excess 
of  ammonia,  the  formation  of  the  white  insoluble  ammonium-magnesium  phosphate 
(p.  467)  serves  as  a  test  also.  Arsenic  acid  (q.v.)  gives  precipitates  of  appearance 
and  composition  similar  to  these  two. 

Phosphorous  Acid.  —  With  cold  water  phosphorus  trioxide  (P4O8) 
yields  phosphorous  acid  very  slowly.  With  hot  water  the  action  is 
exceedingly  violent  and  complex  (p.  464).  This  acid  may  be  obtained 
also  by  the  action  of  water  upon  phosphorus  trichloride,  tribromide, 
or  tri-iodide  and  evaporation  of  the  solution  : 

PC18  +  3H20  ->  P(OH)8  +  3HCL 

A  certain  amount  of  this  acid,  along  with  phosphoric  acid  and  hypo- 
phosphoric  acid,  is  formed  when  moist  phosphorus  oxidizes  in  the  air. 
In  spite  of  the  presence  of  three  hydrogen  atoms,  this  acid  is  dibasic, 
and  two  only  are  replaceable  by  metals.  To  express  this  fact,  the  first 
of  the  following  formulae  is  preferred : 


0 


since  the  symmetrical  formula  would  indicate  no  difference  between  the 
three  hydrogen  atoms.  H  united  directly  to  P,  as  here  and  in  PH8,  is 
not  acidic.  Phosphorous  acid  is  a  powerful  reducing  agent,  precipitat- 
ing silver,  for  example,  in  the  metallic  form  from  solutions  of  its  salts. 
When  heated,  it  decomposes,  giving  the  most  stable  acid  of  phosphorus 
(cf.  pp.  268,  274,  395),  namely,  metaphosphoric  acid,  and  phosphine  : 

4H8P08  -»  3HP08  +  3H20  +  PH8. 

Hypophosphorous  Acid.  —  The  potassium  salt  of  this  acid  is  ob- 
tained, as  we  have  seen,  when  phosphorus  is  heated  with  potassium 
hydroxide  solution  (p.  461).  It  may  be  prepared  in  the  free  form  by 
substituting  barium  hydroxide  for  potassium  hydroxide : 

3Ba(OH)2  +  8P  +  6H,O  -»  3Ba(H2PO,)a  +  2PH8. 

By  careful  addition  of  dilute  sulphuric  acid  to  the  resulting  liquid, 
barium  sulphate  is  precipitated.  On  evaporation  of  the  water  the 
white  crystalline  acid,  H8PO2,  is  obtained.  This  acid  is  monobasic ; 
two  of  its  hydrogen  atoms  cannot  be  displaced  by  metals.  To  ex- 


470  INORGANIC  CHEMISTRY 

f-H 

press  this  fact  the  graphic  formula  0  =  P  1  —  H  is  used.     This  sub- 

(-OH 

stance  is  also  a  powerful  reducing  agent,  tending,  by  the  acquisition  of 
oxygen,  to  pass  into  phosphoric  acid. 

Structural  Formulae  of  Salts  of  Hydrogen.  —  As  a  rule,  the 
formulae  of  acids  have  thus  far  been  written  with  the  ionizable  hydro- 
gen in  front :  HC1,  H2SO4,  HC2H8O2.  This  is  only  one  illustration  of 
the  method  by  which  chemists  have  constantly  sought  to  utilize  form- 
ulae for  the  purpose  of  expressing,  not  merely  the  composition  of  a  sub- 
stance, but  some  of  its  properties  as  well.  By  another  typographical 
device  we  have  attempted  to  indicate  the  behavior  of  dilute  solutions 
by  putting  the  radicals  in  brackets  :  Cu(N03)2,  Ba(OH)2.  These  are 
called  structural  or  constitutional  formulae,  and  their  object  is  not  to 
show  actual  structure,  but  to  exhibit  the  modes  of  action  of  the  sub- 
stance, by  means  of  a  supposed  structure.  Now  the  modes  of  action  of 
a  single  substance  are  often  rather  various,  and  one  and  the  same 
structural  formula  cannot  represent  all  of  these  at  once.  We  have 
observed  this,  particularly,  with  the  oxygen  acids.  Thus,  H2 .  S04  ex- 
presses the  mode  of  activity  in  dilute  solution  and  often  when  no  solvent 
is  present,  as  in  the  action  on  chlorides  (p.  178)  and  nitrates  (p.  439). 
But  when  all  the  hydrogen  of  an  acid  is  not  ionizable,  we  regard  that 
which  is  so  as  part  of  an  hydroxyl  group  in  the  parent  molecules,  and 
the  rest  as  being  attached  to  the  characteristic  non-metal  of  the  acid, 
for  example,  the  phosphorus  (cf.  pp.  391,  392).  Thus,  we  should  write 
phosphorous  acid  POH(OH)2,  instead  of  H2P08H,  to  chronicle  this  fact. 
So  also  the  formula  POH2(OH)  is  used  for  hypophosphorous  acid. 
Molecular  actions,  such  as  those  of  sulphuric  acid  S02(OH)2  (p.  389), 
are  well  shown  by  these  formulae  : 

S02(OH)2  +  2HBr  -*  S02  +  2H2O  +  Br2. 

It  must  be  noted  particularly  that  this  sort  of  formula,  when  the  sub- 
stance for  which  it  stands  is  an  acid,  represents  only  some  features 
in  the  behavior  of  the  anhydrous  substance  and  of  the  molecules,  and 
not  the  ionic  action  in  solution.  A  formula  like  Ba(OH)2,  where  the 
material  is  a  base,  on  the  other  hand,  represents  both  the  ionic  and  the 
molecular  behavior.  The  graphic  formula  is  more  general  (cf.  p.  224). 
It  shows  all  these  relations,  and  often  still  others,  but  none  of  them  so 
specifically. 


PHOSPHORUS  471 

Sulphides  of  Phosphorus. —  Yellow  phosphorus  when  heated  with 
sulphur  unites  with  explosive  violence.  By  using  red  phosphorus  the 
action  can  be  controlled.  By  employing  the  proper  proportions  the 
pentasulphide  P2S5  is  secured.  It  is  purified  by  distillation  from  a 
retort  in  which  a  current  of  carbon  dioxide  is  maintained  (see  below). 
The  distillate  solidifies  to  a  yellow  crystalline  mass,  which  melts  at 
274°  and  boils  at  530°.  Materials  undergoing  chemical  change,  which 
are  to  be  kept  at  a  constant,  high  temperature,  are  often  placed  in 
tubes  suspended  in  the  vapor  of  the  pentasulphide.  When  a  lower 
temperature  is  required,  boiling  sulphur  (448°)  is  used. 

Distillation  in  a  stream  of  some  inactive  gas  is  a  common  means 
of  distilling  under  reduced  pressure  (cf.  p.  276).  The  dilution  of  the 
vapor  lowers  its  partial  pressure,  just  as  would  evacuation.  This  plan 
has  the  advantage,  however,  of  sweeping  the  vapor  away  from  the 
heated  region  into  the  condenser,  and  so  diminishing  the  amount  of 
decomposition.  In  dealing  with  compounds  of  carbon,  a  current  of 
steam  is  often  used  for  the  above  purposes.  It  enables  us  also  to  sepa- 
rate a  slightly  volatile  substance  from  one  which  is  almost  involatile. 

Phosphorus  pentasulphide  acts  upon  water  in  the  cold,  and  upon 
substances  containing  hydroxyl  when  heated  with  them,  the  actions 
being  similar  to  those  of  the  pentachloride  (p.  463)  : 

P2S6  +  8H20  -r>  2HSP04  +  6H,S. 

Other  sulphides,  P4S8,  P2S3,  and  P8S6,  may  be  prepared  by  using  the 
constituents  in  the  proportions  represented  by  these  formulae. 

Comparison  of  Phosphorus  with  Nitrogen  and  with  Sulphur. 

—  Although  phosphorus  and  nitrogen  are  regarded  as  belonging  to  one 
family,  the  differences  between  them  are  more  conspicuous  than  the 
resemblances.  The  latter  are  confined  almost  wholly  to  matters  con- 
cerned with  valence.  The  differences  are  seen  in  the  facts  that  nitro- 
gen is  a  gas  and  exists  in  but  one  form,  while  phosphorus  is  a  solid 
occurring  in  two  varieties,  and  that  the  former  is  inactive  and  the  latter 
active.  The  contrasts  between  phosphine  and  ammonia  (pp.  461-2) 
and  between  the  halides  of  the  two  elements  (p.  462)  have  been  noted 
already.  The  pentoxide  of  nitrogen  decomposes  spontaneously ;  that 
of  phosphorus  is  one  of  the  most  stable  of  compounds.  Nitric  acid 
is  very  active ;  the  phosphoric  acids  are  quite  the  reverse. 

On  the  other  hand,  the  resemblance  of  phosphorus  to  sulphur  is 


472  INORGANIC   CHEMISTRY 

marked.  Both  are  solids,  existing  in  several  forms.  Both  yield  stable 
compounds  with  oxygen  and  chlorine.  The  hydrogen  compounds 
interact  with  salts  to  give  phosphides  of  metals  and  sulphides  of 
metals,  respectively.  Against  these  must  be  set  the  facts,  that  hydro- 
gen sulphide  does  not  unite  with  the  hydrogen  halides  at  all,  and  that 
phosphoric  acid  is  hard  to  reduce,  while  sulphuric  acid  is  reduced  with 
comparative  ease. 

Exercises.  —  1.     Explain  the  effect  of  sulphuric  acid  in  setting 
fire  to  the  earliest  matches  (p.  460). 

2.  Make  a  brief  definition  of  a  substance  capable  of  sublimation 
(p.  463). 

3.  Why  would  a  mixture  of  potassium  dichromate  and  hydrochloric 
acid  (p.  374)  be  less  suitable  than  nitric  acid  for  making  phosphoric 
acid  from  red  phosphorus  ? 

4.  Why  is  not  the  tertiary  phosphate  of  sodium  (p.  467)  decom- 
posed by  heating  ?     What  tertiary  phosphates  would  be  decomposed 
by  this  means  ? 

5.  Formulate  the  hydrolyses  of  the  secondary  and  tertiary  sodium 
orthophosphates  as  was  done  for  sodium  sulphide  (p.  375). 

6.  How  should  you  prepare  Ca2P207  and  Ca(PO3)2  ? 

7.  What   product   should   you   confidently   expect   to   find   after 
heating   (a)  sodium  phosphite,  NaaHP08,  (b)  barium  hypophosphite 
(p.  469)  ? 

8.  Compare  the  elements  chlorine  and  phosphorus  after  the  man- 
ner of  the  comparisons  on  p.  471. 


CHAPTER   XXVIII 

CARBON   AND   THE  OXIDES   OP  CARBON 

The  Chemical  Relations  of  the  Element.  —  The  elements  of  the 
carbon  family  are  carbon,  silicon,  germanium,  tin,  and  lead.  Of  these 
the  first  two  are  entirely  non-metallic,  while  the  others  are  metals 
showing  more  or  less  strong  resemblances  to  the  non-metals.  All  these 
elements  are  quadrivalent  as  regards  the  maximum  valence  which  they 
exhibit.  With  the  exception  of  silicon,  however,  they  all  form  com- 
pounds in  which  they  are  bivalent. 

The  chemistry  of  the  compounds  of  carbon  is  an  exceedingly  ex- 
tensive and  complex  subject.  It  is  commonly  known  as  organic 
chemistry,  on  account  of  the  fact  that  the  majority  of  the  substances 
composing,  and  produced  by,  living  organisms  are  compounds  of  carbon, 
and  that  it  was  at  first  supposed  that  their  artificial  production,  e.g.  with- 
out the  intervention  of  life,  was  impossible.  But  many  natural  organic 
products  have  now  been  made  from  simpler  ones  or  from  the  elements, 
a  process  called  synthesis,  and  the  preparation  of  the  others  is  de- 
layed only  in  consequence  of  difficulties  caused  by  their  instability  and 
complexity.  On  the  other  hand,  hundreds  of  compounds  unknown  to 
animal  or  vegetable  life,  including  many  valuable  drugs  and  dyes,  have 
now  been  added  to  the  catalogue  of  chemical  compounds.  v 

The  elements  entering  into  carbon  compounds  are  chiefly  hydrogen 
and  oxygen.  After  these,  nitrogen,  the  halogens,  and  sulphur  may  be 
named. 

CARBON. 

Occurrence.  —  Large  quantities  of  carbon  are  found  in  the  free 
condition  in  nature.  The  diamond  is  the  purest  natural  carbon,  and 
at  the  same  time  the  least  plentiful.  Graphite,  or  plumbago,  which  is 
the  next  purest,  is  found  in  limited  amounts,  and  is  a  valuable  mineral. 
Coal  occurs  in  numerous  forms  containing  greatly  varying  proportions 
of  free  carbon.  Small  quantities  of  the  free  element  have  been  found 
in  meteorites. 

In  combination,  carbon  is  found  in  marsh-gas,  or  methane  CH4, 
which  is  the  chief  component  of  natural  gas.  The  mineral  oils  consist 

473 


474  INORGANIC   CHEMISTRY 

almost  entirely  of  mixtures  of  various  compounds  of  carbon  and  hydro- 
gen. Whole  geological  formations  are  composed  of  carbonates  of  com- 
mon metals,  particularly  calcium  carbonate  or  limestone,  and  a  double 
carbonate  of  calcium  and  magnesium,  known  as  dolomite. 

The  Diamond.  —  The  varieties  of  carbon  differ  very  markedly  in 
their  physical  properties,  and  to  some  extent  also  in  their  chemical 
behavior.  Diamonds,  which  are  found  in  India,  Borneo,  Brazil,  and 
South  Africa,  are  scattered  sparsely  through  metauiorphic  and  volcanic 
rocks  which  seem  to  have  undergone  secondary  changes.  They  are 
covered  with  a  crust  which  entirely  obscures  their  luster,  and  possess 
natural  crystalline  forms  belonging  to  the  regular  system.  A  form  re- 
lated to  the  octahedron  is  frequently  observed.  It  should  be  noted  that 
this  natural  form  bears  no  relation  whatever  to  the  pseudo-crystal- 
line shape  which  is  conferred  upon 
the  stone  by  the  diamond-cutter. 
Thus,  a  "brilliant"  possesses  one 
rather  large,  flat  face,  which  forms 
the  base  of  a  many  sided  pyra- 
mid (Fig.  91,  showing  two  views). 
This  form  is  given  to  the  stone,  in 
FIG.  91.  order  that  the  maximum  reflection 

of  light   from  its  interior  may  be 

produced.  The  diamond  is  harder  than  any  other  variety  of  matter, 
with  the  exception,  perhaps,  of  one  compound  of  boron,  while  only  one 
or  two  other  materials,  like  carborundum,  approach  it.  It  may  be  re- 
marked in  passing,  that  the  hardness  of  a  substance  is  measured,  on  an 
arbitrary  scale,  by  the  way  in  which  it  is  able  to  scratch  smooth  sur- 
faces of  other  bodies.  The  corner  of  one  of  the  natural  diamond  crys- 
tals will  scratch  the  surface  of  almost  every  other  substance,  while  its 
surfaces  in  turn  are  scratched  by  carbide  of  boron  alone.  Its  specific 
gravity  is  about  3.5,  and  it  is  the  densest  form  of  carbon.  Few  materi- 
als are  capable  of  dissolving  any  of  the  forms  of  carbon.  Molten  iron 
(q.v.)  dissolves  five  or  six  per  cent,  part  of  which  goes  into  combina- 
tion, and  a  few  other  substances  at  high  temperatures  dissolve  much 
smaller  quantities.  The  diamond  is  a  nonconductor  of  electricity. 

The  largest  diamond  known,  the  Jubilee,  was  exhibited  at  the  Paris 
Exposition  of  1900,  and  weighed  49  g.  The  Kohinoor  weighs  22  g. 
The  diamond,  although  its  origin  in  nature  is  still  a  matter  of  uncer- 
tainty, has  been  made  artificially  in  several  ways.  Moissan  (1887) 


CARBON  AND   THE   OXIDES   OP   CARBON  475 

dissolved  carbon  in  molten  iron  and,  after  chilling  the  mass  so  as  to 
produce  a  solid  crust,  which  by  its  shrinkage  severely  compressed  the 
interior,  allowed  the  whole  to  cool  very  slowly.  Portions  of  the 
interior  of  the  ingot  were  treated  with  acid  to  dissolve  the  iron,  and 
amongst  the  insoluble  particles  were  recognized  a  few  microscopic 
fragments  (none  larger  than  0.5  mm.)  which  exhibited  the  form  and 
hardness  of  the  diamond.  The  greater  part  of  the  carbon,  however, 
appeared  as  graphite. 

That  the  diamond  contains  nothing  but  carbon,  is  shown  by  the 
fact  that  when  burned  it  produces  nothing  but  carbon  dioxide. 

Graphite.  —  Graphite  (Gk.  ypa<£<o,  I  write)  is  found  in  Cumber- 
land, Siberia,  Ceylon,  and  elsewhere.  Good  crystals  are  seldom 
found,  but  the  form  appears  to  belong  to  the  hexagonal  system.  The 
mineral  is  extremely  soft,  in  utter  contrast  to  the  diamond,  and  has  a 
smaller  specific  gravity  (about  2.3).  It  conducts  electricity.  It  is  now 
made  artificially  by  an  electro-thermal  process  (cf.  p.  457).  A  power- 
ful alternating  current  is  passed  through  a  mass  of  granular  anthracite, 
and  the  latter,  although  not  melted  by  the  high  temperature,  is  largely 
converted  into  graphite. 

Graphite  is  now  used  exclusively  for  making  the  anodes  in  the 
electrolytic  manufacture  of  chlorine  and  in  related  processes.  Mixed 
with  fine  clay  it  forms  the  "  lead  "  of  lead  pencils.  As  "  black-lead  ", 
it  is  employed  to  protect  ironware  from  rusting.  It  takes  the  place  of 
oil  as  a  lubricant  in  cases  where  the  former  would  be  decomposed  by 
heat.  The  fine,  slippery  scales,  which  it  forms  when  pulverized,  fill 
the  inequalities  in  the  bearings,  and  glide  over  one  another  with  little 
friction. 

Amorphous  Carbon.  —  This  is  the  name  given  to  the  varieties  of 
the  element  which  have  no  crystalline  form.  They  vary  in  specific 
gravity  up  to  1.9.  Coke,  which  is  now  manufactured  in  immense 
quantities  by  heating  coal  until  all  the  volatile  matter  has  been  dis- 
tilled off,  is  a  very  dense  variety  of  amorphous  carbon  used  in  the 
reduction  of  iron  ores  (q.v.}.  If  coal  is  used,  the  draft  is  impeded  by 
softening  of  the  mass.  The  expulsion  of  volatile  matter,  above  the 
zone  of  combustion,  also  causes  waste  of  heat. 

By  the  imperfect  combustion  of  heavy  oils  and  resins,  in  which 
the  flame  plays  upon  a  cooled  surface,  a  finely  divided  form  of  carbon 


476  INORGANIC   CHEMISTRY 

known  as  lampblack  is  produced.  This  is  much  used  in  the  manu- 
facture of  printer's  ink. 

Charcoal  is  chiefly  made  by  the  heating  of  wood  out  of  contact 
with  air.  In  the  more  refined  forms  of  the  process  the  charring  is  con- 
ducted in  retorts,  and  the  materials  which  distil  off  are  used  in  vari- 
ous ways.  Wood  consists  largely  of  cellulose  (C6H1006)n,  incrusted 
with  lignin  and  holding  much  moisture  and  resinous  material.  The 
products  of  its  distillation  are  partly  gaseous  and  partly  fluid.  The 
gases,  consisting  mainly  of  hydrogen,  methane  CH4,  ethane  C2H6,  ethy- 
lene  C2H4,  and  carbon  monoxide  CO,  are  employed,  on  account  of  their 
combustibility,  as  fuel  in  the  distillation  itself.  The  fluids  form  a 
complex  mixture  containing  large  quantities  of  water,  wood  spirit,  or 
methyl  alcohol  CH8OH,  acetic  acid,  acetone  (CH8)2CO,  and  tar.  Wood 
charcoal  exhibits  the  cellular  structure  of  the  material  from  which  it 
was  made,  and  is  therefore  exceedingly  porous.  The  original  mineral 
constituents  of  the  wood  appear  in  the  ash  of  the  charcoal  when  the 
latter  is  burned. 

For  certain  purposes,  charcoals  made  in  the  same  fashion  as  the 
above  from  bones  and  from  blood,  find  wide  application.  The  former, 
called  bone  black  (p.  456),  contains  much  calcium  phosphate.  In  the 
chemical  laboratory,  pure  carbon  is  made,  as  a  rule,  by  the  charring  of 
sugar  (cane-sugar,  C^H^Ojj).  The  sugar  is  purified  from  mineral 
matter,  before  use,  by  crystallization  from  water. 

The  tendency  of  almost  all  carbon  compounds  to  char,  when 
heated,  is  used  as  a  means  of  recognizing  their  presence. 

Properties  of  Charcoal.  —  Charcoal  exhibits  certain  properties 
which  are  not  shared  to  any  extent  by  other  forms  of  carbon.  For 
example,  it  can  take  up  large  quantities  of  many  gases.  Boxwood 
charcoal  will  in  this  way  absorb  ninety  times  its  own  volume  of  am- 
monia, fifty-five  volumes  of  hydrogen  sulphide,  or  nine  volumes  of 
oxygen.  Freshly  made  dogwood  charcoal  (used  in  making  the 
best  gunpowder),  when  pulverized  immediately  after  its  preparation, 
often  catches  fire  spontaneously  on  account  of  the  heat  liberated  by  the 
condensation  of  oxygen.  It  is  therefore  set  aside  for  two  weeks,  tc 
permit  the  slow  absorption  of  moisture  and  air,  before  being  ground  up. 
The  absorbed  gases  may  be  removed  unchanged  by  heating  the  charcoal 
in  a  vacuum.  The  disappearance  of  these  immense  quantities  of  gas 
into  small  pieces  of  charcoal  is  described  as  adsorption,  and  is  caused 
by  the  aefoesion  of  the  gases  to  the  very  extensive  internal  surface 


CARBON  AND  THE  OXIDES  OF  CARBON 


477 


which  the  charcoal  possesses.  Solid  and  liquid  bodies  are  also  in 
many  cases  taken  up  by  charcoal  in  a  similar  fashion.  Thus,  strychnine 
may  be  removed  from  an  aqueous  solution  by  agitation  of  the  latter 
with  charcoal.  In  the  manufacture  of  whiskey  (q.v.)j  the  fusel  oil,  which 
is  extremely  harmful,  is  in  many  cases  removed  by  filtration  of  the 
diluted  spirit  through  charcoal,  before  rectification.  Organic  coloring 
matters,  such  as  litmus  and  indigo,  belong  to  the  class  of  bodies  thus 
extracted  from  solution  by  charcoal.  In  the  refining  of  sugar  the 
syrup  is  boiled  with  charcoal  for  the  purpose  of  removing  a  brown 
resin,  in  order  that  the  product  may  be  perfectly  white.  It  is,  in  part, 
upon  this  property  that  we  rely,  also,  in  the  employment  of  charcoal 
filters.  The  organic  materials  dissolved  in  the  drinking  water  under- 
go adsorption  in  the  charcoal.  In  this  connection,  however,  it  must 
be  remembered  that  the  quantity  which  a  given  mass  of  charcoal  may 
take  up  is  limited,  and  that  careful  cleansing  is  required  in  order  that 
the  efficiency  of  the  filter  may  be  maintained. 

Coal.  —  Peat,  brown  coal,  soft  coal,  and  anthracite  represent,  in  a 
general  way,  different  stages  in  the  decomposition  of  vegetable  matter 
in  absence  of  air.  Water  and  compounds  of  carbon  and  hydrogen  are 
given  off  in  the  process.  The  ratio  which  the  carbon  combined  with 
oxygen  and  hydrogen  bears  to  the  free  carbon  decreases  in  the  order  in 
which  the  substances  stand  above.  The  following  table  shows  this 
change  in  composition  and  the  relations  of  the  substances  to  fresh 
wood  on  the  one  hand  and  charcoal  and  coke  on  the  other : 


Percentage,  excluding  Ash 
and  Moisture,  of  : 

Percentage 
Ash. 

Per- 
centage of 
Water 
(Air  Dried). 

Calorific 
Value 

Calories. 

C 

H 

0 

N 

Wood    .      . 

45 

6 

48 

1 

1.6 

18-20 

2700 

Peat      .     . 

60 

6 

32 

2 

6—20 

20-30 

3500 

Brown  coal 

70 

5 

24 

1 

3—30 

16 

30-6000 

Soft  coal     . 

82 

6 

12 

1 

1—16 

4 

66-8000 

Anthracite 

94 

3 

3 

.  , 

1.6 

2 

70-8000 

Charcoal     . 

96 

1.7 

3.4 

4 

6.5 

8080 

Coke      .      . 

96 

0.7 

2.6 

1 

3.4—11 

2 

7700 

Chemical  Properties  of  Carbon.  —  Diamond,  graphite,  and 
amorphous  carbon  probably  differ  from  one  another,  not  merely  in 
physical  properties,  but  also  chemically.  Certainly  the  stability  of 


478  INORGANIC   CHEMISTRY 

compounds  containing  many  units  of  carbon  in  their  molecules  indi- 
cates a  great  tendency  of  carbon  to  combine  with  itself,  and  gives 
plausibility  to  the  belief  that  the  molecule  of  free  carbon  may  itself  be 
complex.  Differences  in  the  size  or  structure  of  these  complex  mole- 
cules would  account  for  the  variety  in  the  forms  of  the  element. 
Amorphous  carbon  is  the  least  stable  of  the  three,  for  it  liberates  most 
heat  in  entering  into  combination.  Since  graphite  is  formed  at  high 
temperatures,  and  diamonds  turn  into  a  black  mass  under  the  same 
conditions,  we  may  presume  that  graphite  is  the  most  stable,  at  least 
at  3000°. 

The  most  common  uses  of  carbon  depend  upon  its  great  tendency 
to  unite  with  oxygen,  forming  carbon  dioxide.  Under  some  circum- 
stances carbon  monoxide  is  produced.  Aside  from  the  direct  employ- 
ment of  this  action  for  the  sake  of  the  heat  which  is  liberated,  it  is 
used  also  in  the  reduction  of  ores  of  iron,  copper,  zinc,  and  many  other 
metals.  When,  for  example,  finely  powdered  cupric  oxide  and  carbon 
are  heated,  copper  is  obtained.  The  gas  given  off  is  either  carbon 
dioxide,  or  a  mixture  of  this  with  carbon  monoxide,  according  to  the 
proportion  of  carbon  used  : 

CuO  +  C  ->    Cu  -f  CO, 
2CuO  -f-  C  — >  2Cu  +  C02. 

Carbon  unites  directly  with  hydrogen  very  reluctantly.  When  an 
electric  arc  is  produced  between  carbon  poles  in  a  tube  through  which 
a  stream  of  hydrogen  passes,  acetylene  C2H2  is  formed.  The  presence 
of  this  gas  may  be  shown  by  the  luminosity  its  combustion  confers 
on  the  hydrogen  flame.  This  substance  can  form  the  starting-point 
for  the  artificial  preparation  of  many  carbon  compounds,  and  its  syn- 
thesis possesses  therefore  a  certain  interest. 

At  the  high  temperatures  produced  in  the  electric  furnace,  carbon 
unites  with  many  metals  and  some  non-metals.  Compounds  formed 
in  this  way  are  known  as  carbides,  such  as  aluminium  carbide  A14C8, 
calcium  carbide  CaC2,  and  carborundum  CSi. 

Calcium  Carbide.  —  This  compound,  which  is  colorless  when 
pure,  is  manufactured  in  an  electric  furnace,  by  the  interaction  of  finely 
pulverized  limestone  or  quicklime  with  coke  : 

CaO  +  3C  — >  CaC2  -f  CO. 

The  operation  is  a  continuous  one,  the  materials  being  thrown  into  the 
left  side* of  the  drum  (Fig.  92,  diagrammatic),  and   the   product   re- 


CARBON  AND  THE  OXIDES  OF  CARBON 


479 


moved  on  the  right.    The  carbon  poles  are  fixed.      The  arc  having 
been  established,  the  drum  is  rotated  slowly  as  the  carbide  accumu- 
lates.    The  current  enters  by  one  carbon,  passes  through  the  carbide, 
and  leaves  by  the  other.     The  high  resistance  of  the  partially  trans- 
formed material  causes  the  production  of  the  heat.     When  the  action 
in  one  layer  approaches  completion,  the  resistance  falls,  the  current 
increases,  and  an  armature  round  which  the  wire  passes  (not  shown  in 
Fig.  92)  comes  into  operation  and  turns  the  drum.     In  this  way  the 
carbide  just  formed  is  continuously  moved  away  from  the  carbons,  and 
new  material,  introduced  on  the  left,  falls  into  the  path  of  the  current. 
The  iron  plates  which  form  the  circumference  of  the  drum  are  added 
on  the  left  and  removed  on  the 
right,  where  also  the  carbide  is 
broken  out  with  a  chisel.     The 
drum   revolves   once   in  about 
three   days.      The   product   is 
used  for  making  acetylene  (q.v.). 

CARBON  DIOXIDE  AND 
CARBONIC  ACID. 

Occurrence.  —  Carbon  di- 
oxide is  present  in  the  atmos- 
phere, and  issues  from  the 
ground  in  large  quantities  in 
certain  neighborhoods,  as,  for 
example,  near  the  Lake  of 
Laach,  in  the  so-called  Valley  of 

Death  in  Java,  and  in  the  G-rotta  del  Cane  near  Naples.  Effervescent 
mineral  waters  contain  it  in  solution,  and  their  effervescence  is  caused 
by  the  escape  of  the  gas  when  the  pressure  is  reduced.  Well-known 
waters  of  this  kind  are  those  of  Selters  (whence,  by  a  singular  per- 
version, the  English  word  seltzer  is  derived)  and  of  the  Geyser 
Spring  at  Saratoga. 

Modes  of  Formation.  —  Carbon  dioxide  is  produced  by  combus- 
tion of  carbon  in  the  presence  of  an  excess  of  oxygen  : 

C  +  O  ->  C0. 


FIG.  92. 


The  combustion  of  all  compounds  of  carbon,  as  well  as  the  slow  oxida- 


480  INORGANIC   CHEMISTRY 

tion  in  the  tissues  of  plants  and  animals,  leads  to  the  formation  of  the 
same  product. 

It  was  Joseph  Black  (1757)  who  first  recognized  the  gas  as  a 
distinct  substance.  He  observed  its  formation  when  marble  or  mag- 
nesium carbonate  was  heated : 

CaC03?=»CaO  -f  CO2, 

and  named  the  gas  "  fixed  air  "  from  the  fact  that  it  was  contained  in 
these  solids.  The  above  action  had  been  used  for  centuries  in  making 
quicklime  (calcium  oxide).  All  common  carbonates,  excepting  the 
normal  carbonates  of  potassium  and  sodium,  decompose  in  this  way, 
leaving  the  oxide  of  the  metal. 

Black  found  that  the  gas  was  also  produced  when  acids  acted  upon 
carbonates,  and  this  method  is  commonly  employed  in  the  laboratory  : 

CaC08  (solid)  *=>  CaCOs  (diss'd)  ^  Ca"+C0,"  > 

2HC1  (diss'd)  <=>  2C1'  +  2H*  J  ^  H*CO*^  H*° 

Since  the  carbonic  acid  is  very  slightly  ionized,  the  action  is  like 
that  of  acids  on  sulphites  (p.  398).  The  carbonate  of  calcium,  how- 
ever, is  very  slightly  soluble,  so  that  an  additional  equilibrium  controls 
its  solution.  In  this  respect  the  action  is  like  that  of  acids  on  ferrous 
sulphide  (p.  371). 

Carbon  dioxide  is  also  a  product  of  the  fermentation  of  sugar  (q.v.), 
as  Black  had  the  credit  of  showing. 

Physical  Properties.  —  Carbon  dioxide  is  a  colorless,  odorless 
gas.  It  is  one-half  heavier  than  air.  The  G.M.V.  weighs  44  g.  The 
critical  temperature  is  31.1°.  Liquefied  carbon  dioxide  boils  at  —79°. 
The  sp.  gr.  of  the  liquid  at  0°  is  0.95.  At  0°  its  vapor  tension  is 
35.4  atmospheres  and  at  20°  59  atmospheres.  It  must  be  preserved, 
therefore,  in  very  strong,  wrought-iron  cylinders.  Large  quantities  of  it, 
often  collected  from  fermentation  vats,  are  sold  in  such  cylinders,  and 
used  in  operating  beer-pumps  and  in  making  aerated  waters.  When 
the  liquid  is  allowed  to  flow  out  into  an  open  vessel  it  cools  itself  by 
its  own  evaporation  and  forms  a  white,  snowlike  mass.  Solid  carbon 
dioxide  evaporates  without  melting  (cf.  p.  463).  A  mixture  of  solid 
carbon  dioxide  with  ether  is  frequently  used  as  a  freezing  mixture 
(—80°).  The  ether  is  employed  to  secure  better  contact  with  the  body 
to  be  cooled. 


CARBON   AND   THE   OXIDES   OF   CARBON  481 

The  great  contrast  in  the  speeds  of  a  chemical  change  at  two 
cemperatures  (cf.  p.  72)  may  be  illustrated  by  putting  a  minute  piece 
of  sodium  in  some  30  per  cent  hydrochloric  acid  which  has  been 
cooled  in  the  above  mixture.  Hardly  any  interaction  can  be  observed. 
But  if  the  temperature  of  the  acid  is  allowed  to  rise,  the  action  becomes 
more  and  more  rapid,  and  ends  by  being  explosively  violent. 

Carbon  dioxide  gas  under  a  pressure  of  7.60  mm.  and  at  a  tempera- 
ture of  15°  dissolves  in  its  own  volume  of  water.  Up  to  pressures 
of  four  or  five  atmospheres  Henry's  law  describes  its  solubility  accu- 
rately. An  aqueous  solution,  prepared  under  a  pressure  of  2-3  atmos- 
pheres, is  familiarly  known  as  soda  water. 

Chemical  Properties.  —  Carbon  dioxide  is  a  stable  compound. 
At  2000°  (760  mm.  press.)  about  7.5  per*cent  is  dissociated:  2CO2  ^± 
2CO  -j-  02.  Water  is  decomposed  only  one-fourth  as  much. 

The  most  active  metals,  such  as  potassium,  sodium,  and  magnesium, 
burn  brilliantly  when  heated  in  carbon  dioxide,  producing  the  oxide 
of  the  metal  and  free  carbon.  Less  active  metals,  such  as  zinc  and 
iron,  give  an  oxide  of  the  metal  and  carbon  monoxide  (q.v.). 

Carbon  dioxide  unites  directly  with  many  oxides,  particularly  those 
of  the  more  active  metals,  such  as  the  oxides  of  potassium,  sodium, 
calcium,  etc.  Hence  the  decomposition  of  calcium  carbonate  by  heat- 
ing (p.  480)  is  a  reversible  action,  which  proceeds  in  the  opposite  direc- 
tion when  a  sufficient  pressure  of  carbon  dioxide  is  employed  (cf.  p. 
256  and  Chap.  xxxv). 

Carbon  dioxide,  when  dissolved  in  water,  forms  an  unstable  acid  : 


The  name  carbonic  acid  is  frequently,  though  improperly,  given  to 
the  gas  itself,  which  is  really  the  anhydride  of  the  acid  and  has  no 
acid  properties. 

Chemical  Properties  of  Carbonic  Add.  —  The  solution  of  car- 
bon dioxide  in  water  exhibits  the  properties  of  a  weak  acid.  It  con- 
ducts electricity,  although  not  well.  It  turns  litmus  red,  though  not 
so  decidedly  as  do  strong  acids.  Its  feebleness  is  due,  however,  not 
exclusively  to  the  small  degree  of  ionization,  but  also  to  the  fact  that 
ordinary  solutions  of  carbon  dioxide  are  necessarily  very  dilute.  The 
ionization  takes  place  chiefly  according  to  the  equation  : 

H'  -f  HC08'. 


482  INORGANIC   CHEMISTRY 

In  a  deci-norrnal  solution,  less  than  two  molecules  of  the  acid  in  a 
thousand  are  ionized.  The  conditions  of  equilibrium  between  the  gas 
and  the  solution  are  precisely  similar  to  those  described  under  sul- 
phurous acid  (p.  393). 

Carbonates.  —  When  excess  of  an  aqueous  solution  of  carbonic 
acid  is  mixed  with  a  solution  of  a  base  like  sodium  hydroxide,  or,  as 
the  operation  is  more  usually  performed,  when  carbon  dioxide  is 
passed  directly  into  a  solution  of  the  alkali,  water  is  formed  and  the 
carbonate  remains  dissolved: 

H2CO3  +  NaOH  ±5  H2O  +  NaHC03,  or  H'  +  OH'  ->  H20. 

The  product  is  sodium  hydrogen  carbonate  (sodium  bicarbonate). 
Although  technically  an  acid  salt,  its  solution  is  neutral  on  account  of 
the  exceedingly  slight  dissociation  of  the  HCO/  ion.  By  addition  of 
an  equivalent  of  sodium  hydroxide  to  the  solution  of  the  bicarbonate 
the  normal  carbonate  is  obtained  : 

NaOH  +  NaHC03  +±  H20  +  Na^CO,,. 

This  solution  is  alkaline  in  reaction,  for  the  same  reason  that  a  solution 
of  secondary  sodium  orthophosphate  is  so  (cf.  p.  466). 

The  carbonates,  with  the  exception  of  those  of  potassium,  sodium, 
and  ammonium,  are  insoluble  in  water,  and  may  be  obtained  by  precipi- 
tation when  the  proper  ions  are  employed.  For  example  : 

MgSO<  +  Na2C08  <=>  MgC08  J+  Na2S04  or   Mg"+  CO/'  <=±  MgC03  J. 

The  aqueous  solution  of  carbon  dioxide  interacts  with  solutions  of 
barium  and  calcium  hydroxides  in  a  similar  manner  : 

Ca(OH)2  +  H2CO8 fc?  CaC08  |  +  H2O. 

The  formation  of  precipitates  with  these  solutions  is  used  as  a  test  for 
carbon  dioxide  and  a  means  of  estimating  its  amount  in  a  sample  of 
air  (q.v.). 

Excess  of  carbon  dioxide  converts  calcium  carbonate  into  the  more 
soluble  bicarbonate,  and  hence  considerable  quantities  of  "  lime  "  are 
frequently  held  in  solution  by  natural  waters : 

I^COg  +  CaCO8  <±  H2Ca(C08)2. 

A  considerable  excess  of  carbon  dioxide  is  required  to  convert  the 
whole  of  the  carbonate  into  the  soluble  bicarbonate,  since  the  action 


CARBON   AND   THE   OXIDES   OF   CARBON  483 

is  markedly  reversible.  In  the  same  fashion,  the  carbonates  of  iron 
(FeC03),  magnesium,  and  zinc  are  somewhat  soluble  in  water  con- 
taining free  carbonic  acid.  In  fact,  the  solution,  transportation,  and 
deposition  of  all  these  carbonates  take  place  in  nature  on  a  large  scale 
by  the  alternate  progress  and  reversal  of  this  action.  Water  contain- 
ing calcium  carbonate  in  solution  is  known  as  hard  water  (g.v.),  that 
containing  ferrous  carbonate  as  chalybeate,  or  iron  water. 

Mole  of  Chlorophyll-bearing  Plants  in  Storing  Energy.— 

While  in  plants  the  same  consumption  of  oxygen  and  production  of 
carbon  dioxide  goes  on  as  in  animals,  only  with  less  rapidity,  an  action 
which  is  in  a  general  way  a  reversal  of  this  takes  place  at  the  same 
time.  The  chlorophyll  and  protoplasm  in  the  leaves  of  the  plant  have 
the  power  of  taking  up  carbon  dioxide.  Part  of  the  oxygen  is  restored 
to  the  air,  and  the  rest  of  the  substance,  including  all  the  carbon,  is 
used  by  the  plant  as  food.  This  operation  goes  on  only  in  sunlight 
(see  below).  The  details  of  the  chemical  changes  are  not  thoroughly 
understood,  but  the  various  chemical  compounds  which  plants  con- 
struct in  large  quantities,  of  which,  sugar,  starch,  and  cellulose  are 
prominent  examples,  are  built  up  as  the  result  of  this  action.  In  a 
rough  fashion,  and  disregarding  the  steps  by  which  the  process  takes 
place,  we  may  represent  the  chemical  change  by  means  of  the  thermo- 
chemical  equation : 

6C02  +  5H20  ->  C6H10O5  -f  602  —  671,000  calories. 

Since  the  production  of  carbon  dioxide  by  the  combustion  of  any 
organic  compound  gives  out  heat,  the  partial  reversal  of  this  combus- 
tion, of  which  the  green  parts  of  the  plant  are  the  scene,  requires  the 
expenditure  of  energy,  and  the  source  of  this  energy  is  to  be  found  in  the 
sunlight.  The  difference  in  total  energy  between  water  and  carbon 
dioxide,  on  the  one  hand,  and  the  cellulose,  starch,  or  sugar  and  free 
oxygen  on  the  other,  is  very  considerable.  The  above  figures  indicate 
roughly  (p.  79)  this  difference  (=  the  amount  of  energy  stored)  for 
cellulose,  and  the  values  for  the  other  compounds  are  of  the  same 
order. 

The  importance  of  this  remarkable  endothermal  action,  involving 
the  storing  of  the  energy  of  sunlight,  is  very  great.  Aside  from  a 
little  work  done  by  water-power,  the  whole  energy  used  by  man  and 
by  animals  comes  from  the  reversal  of  it.  The  compounds  forming 
the  structure  of  the  plant  are  employed  in  several  different  ways  for 


484  INORGANIC   CHEMISTRY 

this  purpose.  When  consumed  by  herbivora  as  food,  the  chemical 
changes  which  they  undergo  furnish  the  energy  necessary  for  the  con- 
tinued life  of  the  organism.  The  whole  of  the  material  is  not  at  once 
reduced  to  the  state  of  carbon  dioxide,  but  passes  into  other  forms  of 
combination,  which  in  turn  become  the  food  of  the  carnivora,  or  flesh- 
eating  animals.  By  the  oxidation  to  carbon  dioxide  which  takes  place 
in  the  bodies  of  these  animals,  the  process  of  exhausting  the  possible 
energy  of  the  carbon  compounds  is  completed. 

In  another  fashion  we  secure -energy  from  the  materials  of  plants 
by  burning  wood  and  employing  the  heat  thus  produced.  In  still 
another  fashion,  after  the  wood  has  undergone  partial  decay  and 
conversion  into  coal,  we  secure  the  remaining  energy  which  this  con- 
tains by  its  final  combustion  in  the  furnace  of  the  steam-engine. 

It  should  be  noted  that  the  energy  in  the  last  case  is  not  stored  exclusively  in  the 
coal,  but  is  shared  between  carbon  and  the  oxygen  of  the  air.  If  our  atmosphere 
consisted  of  compounds  of  carbon,  then  the  material  corresponding  to  stores  of  coal 
would  have  to  be  oxygen  or  compounds  of  oxygen,  and  we  should  be  likely  then 
to  speak  of  the  energy  as  being  stored  for  us  and  sold  in  the  form  of  oxygen.  That 
we  are  in  the  habit  of  speaking  of  it,  at  present,  as  going  with  the  carbon  is  because 
the  oxygen  of  the  air  is  supplied  free  of  charge,  while  the  coal  and  wood  have 
to  be  purchased. 


.  Photochemical  Action.  —  We  have  seen  that  light  may  simply 
act  catalytically,  as  on  a  mixture  of  hydrogen  and  chlorine  (p.  174)  or 
an  aqueous  solution  of  hypochlorous  acid  (p.  268).  These  actions  in- 
volve the  liberation  of  energy  and  go  on  spontaneously  (cf.  p.  271)  under 
proper  conditions.  On  the  other  hand,  light  may  actually  be  consumed 
in  large  amount  in  producing  a  chemical  change,  as  in  decomposing 
silver  chloride  (p.  14),  or  in  the  above  instance.  All  wave  lengths 
of  light,  which  is  the  same  as  to  say  all  colors  of  light,  are  not  equally 
active  in  any  one  case.  But  there  is  no  particular  set  of  wave  lengths 
which  is  of  special  chemical  activity.  In  the  action  on  silver  chloride, 
green  and  blue  light  is  very  active,  while  red  is  almost  without  effect. 
Here,  in  the  actions  in  which  chlorophyl  is  concerned,  it  is  the  red 
and  yellow  light  that  produces  the  chemical  change,  and  a  plant  ex- 
posed to  blue  light  (e.g.,  by  shading  with  blue  glass)  will  assimilate 
none  of  the  carbon  dioxide  in  the  air  surrounding  it.  The  chemical 
substances  in  the  retina  of  the  eye  seem  to  resemble  those  in  the  leaves 
of  plants,  for  they  are  most  affected  by  red  and  yellow  light.  To  put 
this  another  way,  a  spectrum  of  uniform  intensity  throughout,  when 


CARBON   AND   THE   OXIDES   OF   CARBON  485 

viewed  by  a  plant  or  a  human  eye,  would  appear  to  be  brightest  in 
the  red  and  yellow  portions,  while  a  considerable  stretch  towards  the 
blue  extremity  would  actually  be  invisible.  On  the  other  hand,  to  an 
eye  in  which  the  active  substance  was  silver  chloride,  if  such  an  eye 
could  be  imagined  as  existing,  the  red  end  would  be  invisible  and  the 
blue  and  ultra-violet  would  be  the  most  brilliant  parts. 

CAKBON  MONOXIDE. 

Preparation.  —  Carbon  monoxide  is  formed  in  many  industrial 
operations.  We  commonly  observe  the  blue  flame  of  burning  carbon 
monoxide  playing  on  the  surface  of  a  coal  fire.  The  gas  is  produced 
by  the  passage  of  the  carbon  dioxide,  which  is  first  formed,  through  the 
upper  layers  of  heated  coal : 

C02  4-  C  — >  2CO. 

A  similar  reduction  of  carbon  dioxide  is  produced  by  metals  such  as 
zinc,  when  a  moderate  heat  is  applied : 

C02  -f-  Zn  -»  ZnO  +  CO. 

On  a  large  scale,  a  mixture  of  carbon  monoxide  and  hydrogen  is 
prepared  as  the  basis  of  -water  gas.  Steam  is  turned  into  an  iron 
cylinder  lined  with  fire  clay  and  filled  with  vigorously  burning  coke : 

C  -f  H.O  ->  CO  +  H,. 

The  products  are  both  combustible,  and,  by  the  addition  of  substances 
which  burn  with  a  luminous  flame,  the  mixture  is  used  for  the  manu- 
facture of  illuminating-gas  (q.v.). 

In  the  laboratory,  carbon  monoxide  is  frequently  obtained  by  heat- 
ing oxalic  acid,  a  solid,  white,  crystalline  substance,  with  concentrated 
sulphuric  acid.  The  latter  is  here  employed  simply  as  a  dehydrating 
agent  (p.  388),  so  that  it  need  not  be  included  in  the  equation  : 

H2C204  ->  C02  +  CO  +  H20. 

To  obtain  pure  carbon  monoxide  from  this  mixture  it  is  necessary  to 
remove  the  carbon  dioxide  and  to  dry  the  gas.  The  carbon  dioxide 
may  be  absorbed  by  passing  the  gas  through  a  concentrated  solution 
of  potassium  hydroxide.  By  treatment  of  formic  acid,  or  sodium  for- 
mate, with  sulphuric  acid,  the  presence  of  the  carbon  dioxide  may  be 
avoided : 

HCH02  ->  CO 


486  INORGANIC   CHEMISTRY 

Physical  Properties.  —  Carbon  monoxide  is  a  colorless,  tasteless, 
odorless  gas.  It  is  very  slightly  soluble  in  water.  Its  density  is 
almost  the  same  as  that  of  air,  for  the  G.M.V.  weighs  28  g.  When 
liquefied  it  boils  at  -190°. 

Chemical  Properties.  —  All  the  chemical  properties  of  carbon 
monoxide  are  referable  to  the  fact  that  in  it  the  element  carbon 
appears  to  be  bivalent :  C  =  O.  The  compound  is  in  fact  unsaturated, 
and  combines  with  oxygen,  chltirine,  and  other  substances  directly. 
Thus  the  gas  burns  in  the  air,  uniting  with  oxygen  to  form  carbon 
dioxide.  Again,  iron  (q.v.)  is  manufactured  by  the  reduction  of  the 
oxide  of  iron  by  gaseous  carbon  monoxide  in  the  blast  furnace : 

Fe208  +  3CO  <=»  2Fe  +  3CO2. 

In  sunlight  carbon  monoxide  unites  directly  with  chlorine  to  form 
carbonyl  chloride  COC12.  It  is  absorbed  by  a  solution  of  cuprous 
chloride  in  hydrochloric  acid  or  ammonium  hydroxide,  forming  a  com- 
pound whose  composition  is  probably  represented  by  the  formula 
Cu2Cl2,CO,2H20.  It  unites  directly  with  certain  metals,  notably  nickel 
and  iron,  with  which  it  forms  the  so-called  nickel  carbonyl  and  iron  car- 
bonyl, respectively.  The  former  is  a  colorless,  volatile  liquid  Ni(CO)4. 

The  gas  is  an  active  poison.  When  inhaled  it  unites  with  the 
haemoglobin  of  the  blood  to  the  exclusion  of  the  oxygen,  which  forms 
with  the  haemoglobin  a  less  stable  compound  (cf.  p.  432).  A  quan- 
tity equivalent  to  about  10  c.c.  of  the  gas  per  kilo,  weight  of  the 
animal  is  sufficient  to  produce  death,  about  one-third  of  the  whole 
haemoglobin  having  entered  permanently  into  combination  with  carbon 
monoxide. 

The  quantities  of  heat  given  out  by  the  successive  unions  of  two  units  of  oxy- 
gen with  one  unit  of  carbon  are  worth  recording  : 

C  +  O    — >  CO  +  29,650  calories, 
CO  +  O  — >  C02+  68,000  calories. 

It  will  be  seen  that  the  addition  of  the  second  atom  of  oxygen  appears  to  cause  the 
evolution  of  a  very  much  larger  amount  of  heat  than  does  that  of  the  first.  It 
must  be  remembered,  however,  that  the  carbon  monoxide  is  gaseous,  while  the 
carbon  in  the  first  equation  is  solid,  and  probably  in  a  condition  of  complex  molec- 
ular aggregation.  The  heats  produced  by  the  unions  of  the  two  units  are  probably 
not  very  different,  but  in  the  first  case  a  large  amount  of  the  heat  is  used  up  in 
disintegrating  the  carbon  and  bringing  it  iuto  the  gaseous  condition. 


CARBON   AND   THE   OXIDES  OF   CARBON  487 

CARBONYL  CHLORIDE  AND  UREA. 

Carbonyl  Chloride.  —  This  substance  is  also  named  phosgene 
(Gk.  <£<os,  light  ;  yewai/,  to  produce),  on  account  of  its  formation  by  the 
catalytic  influence  of  sunlight  (p.  484).  On  a  commercial  scale  it  is 
obtained  by  passing  the  mixed  carbon  monoxide  and  chlorine  over 
animal  charcoal,  which  assists  the  union  catalytically.  It  is  a  liquid 
which  boils  at  8°,  possesses  a  suffocating  odor,  and  is  very  soluble  in 
benzene  and  some  other  hydrocarbons.  When  brought  into  contact 
with  water  it  is  hydrolyzed  at  once,  forming  carbonic  acid  and  hydro- 
chloric acid  : 

COCLj  +  2H20  -^CO,  +  2HC1. 


Urea.  —  When  ammonia  and  carbonyl  chloride  are  mixed  in  the 
proper  proportions,  urea,  a  most  interesting  chemical  substance,  is 
produced  : 

Cl         H-NH2  NH, 

0  =  C/        +  -»0  =  C'          +  2HC1 

XC1         H-NH2  XNH2 

Excess  of  ammonia  has  to  be  used  to  combine  with  the  hydrogen 
chloride  thus  set  free,  so  that  the  final  equation  is  : 

COC12  +  4NH3  -»  CO(NH2)2  +  2NH4C1. 

The  urea,  a  white,  crystalline  solid,  is  soluble  in  alcohol,  while  ammo 
nium  chloride  is  not,  so  that  the  former  may  be  washed  out  by  means  of 
this  solvent  and  recovered  by  evaporation.  A  little  reflection  will 
show  that,  using  the  above  action  as  the  final  stage,  urea  can  be  built 
up  from  the  simple  substances  composing  it. 

Urea  was  known  long  before  any  method  for  its  synthesis  had 
been  discovered.  It  is  the  chief  product  of  the  decomposition  of  com- 
pounds of  nitrogen  in  the  animal  body,  and  is  found  in  the  liquid 
excrements  of  animals.  It  was  regarded  as  a  typical  organic  substance, 
in  the  old  sense  of  the  word  (p.  473).  In  1828  Wohler  succeeded  in 
preparing  it  artificially  (see  below).  This  was  the  first  synthesis  by  a 
chemist,  of  a  true  "  organic  "  substance,  and  its  preparation  proved  to 
be  the  precursor  of  many  discoveries  of  a  similar  nature.  From  a 
later  year,  about  1840,  we  may  date  the  transition  of  organic  chem- 
istry, a  science  in  which  the  mystery  of  life  was  supposed  to  be 
supreme,  into  the  chemistry  of  the  compounds  of  carbon,  which  is  a 
branch  of  inorganic  chemistry. 


488  INORGANIC   CHEMISTRY 

WOhler  used  ammonium  cyanate  (g.fl.),  a  substance  in  whose  preparation  we 
are  independent  of  all  products  of  life  processes.  When  ammonium  cyanate,  or  a 
mixture  of  any  ammonium  salt  with  potassium  cyanate  in  solution  in  water,  is 
warmed  for  some  time,  an  intramolecular  change  (cf.  p.  15)  takes  place,  and  long 
prisms  of  urea  are  deposited  as  the  liquid  cools  : 

NH4.CNO  <=±  CO(NH2)2. 

Since  the  action  is  reversible,  about  four  or  five  per  cent  of  the  ammonium  cyanate 
remains  unchanged. 

The  two  substances  just  mentioned  are  entirely  different  in  chemical  proper- 
ties. Ammonium  cyanate  is  a  highly  ionized  salt,  while  urea  is  not  a  salt  at  all, 
but  a  substance  like  ammonia  which  unites  with  acids  to  form  salts.  Materials 
which,  like  these,  have  the  same  composition  and  the  same  numbers  of  units  in 
their  molecules,  and  yet  possess  different  properties,  are  spoken  of  in  chemistry  as 
isomera.  The  formulae  we  have  employed  attempt  to  explain  the  differences  in 
their  properties  by  suggesting  a  difference  in  their  molecular  structure  (cf.  p.  224). 

Assisted  by  the  catalytic  action  of  certain  ferments,  urea,  when  dis- 
solved in  water,  can  take  up  two  molecules  of  the  solvent  to  form 
ammonium  carbonate  : 


CO(NH2)2  +  2HP  ->  (NH4)2C08  <=±  2NH8  +  H.O  +  C02. 

Ammonium  carbonate  (q.v.)  is  a  somewhat  unstable  compound,  and  in 
turn,  gives  off  ammonia  and  carbon  dioxide.  To  this  action  is  due  in 
part  the  pronounced  odor  of  ammonia  arising  from  the  decomposition 
of  sewage. 

Carbon  Disulphide.  —  This  compound  is  used  in  inorganic  chem- 
istry chiefly  as  a  solvent.  It  is  made  by  direct  union  of  sulphur 
vapor  and  glowing  charcoal.  An  electro-thermal  method  of  carrying 
this  out  employs  a  furnace  like  that  in  Fig.  90  (p.  457).  The  sub- 
stance comes  off  as  a  vapor  and  is  condensed. 

Carbon  disulphide  is  a  colorless,  highly  refracting  liquid.  When 
pure  it  possesses  a  pleasant  odor,  but  traces  of  other  compounds  give 
the  commercial  article  a  disagreeable  smell.  It  boils  at  46°  and  burns 
in  air,  forming  carbon  dioxide  and  sulphur  dioxide.  Iodine,  phos- 
phorus, sulphur,  rubber,  and  other  substances  dissolve  freely  in  it. 

Exercises.  —  1.  To  which  of  the  factors  in  the  interaction  of  cal- 
cium carbonate  and  hydrochloric  acid  (p.  480)  is  due  the  forward  dis- 
placement of  all  the  equilibria  ? 

2.   What  will  be  the  excess  of   pressure  inside  a  bottle  of  soda- 


CARBON  AND  THE   OXIDES  OF  CARBON  489 

water  when  four  volumes  of  carbon  dioxide   are   dissolved   in   one 
volume  of  water  ? 

3.  What  volume  of  liquid  carbon  dioxide,  measured  at  0°,  will  be 
required  to  give  75  liters  of  the  gas  at  0°  and  760  mm.  pressure  ? 

4.  What  will  be  the  effect  of  increase  in  pressure  on  the  dissocia- 
tion of  carbon  dioxide  (p.  481)  ? 

5.  Prepare  a  diagram  showing  the  whole  scheme  of  equilibria  in- 
volved in  the  hydrolysis  of  sodium  carbonate  (p.  482). 

6.  What  volume  of  carbon  dioxide  at  0°  and  760  mm.  is  required 
for  complete  interaction  with  one  liter  of  normal  sodium  hydroxide  ? 


CHAPTER   XXIX 
SOME   CARBON   COMPOUNDS 

THE  compounds  of  carbon  with  hydrogen  are  called  hydrocarbons. 
Those  containing  oxygen  as  well  are  divided  into  numerous  and 
extensive  groups  according  to  their  behavior.  Thus  there  are  acids 
like  acetic  acid,  carbohydrates  like  sugar  and  starch,  alcohols  like 
common  (ethyl)  alcohol,  esters  like  ethyl  acetate  and  fat,  ethers  like 
common  (ethyl)  ether.  There  are  also  bodies  related  to  cyanogen, 
like  prussic  acid,  which  contain  nitrogen.  We  can  discuss  only  one 
or  two  examples  from  each  of  the  groups  named. 

THE   HYDROCARBONS. 

More  than  two  hundred  and  fifty  compounds  of  carbon  and  hydro- 
gen have  been  described.  They  fall  into  several  distinct  series,  the 
chief  one  of  which  contains  methane  CH4  as  its  simplest  member. 
On  account  of  the  fact  that  certain  members  of  this  set  are  found  in 
paraffin,  it  is  commonly  known  as  the  paraffin  series.  For  the  reason 
that  in  this  series  the  carbon  has  all  its  four  valences  employed,  the 
members  are  also  called  the  saturated  hydrocarbons. 

Paraffin  Series  of  Hydrocarbons.  —  The  following  list  gives 
the  formulae  of  a  few  members  of  this  series,  with  their  names  and 
their  boiling-points  or  melting-points : 

CH4  .      .      .  methane b.-p.     -  164°    -N 


ethane 


C3H8.  .  propane  '  °"°       > Gases 


butane 


pentane 
hexane 


-89.5° 


-37 


J 


35C 
71C 


heptane     ...  99°       Liquids 


490 


SOME   CARBON   COMPOUNDS  491 

i«H34     •      •     hexadecane m.-p.        18°      "1  b.-p.  287.5° 

H72     .      .     pentatriacontane     .     .      .     "  74.7°  I- Solids 


.      .     hexacontane  (dimyricyl)  .     "  102° 

It  will  be  noted  that  the  hydrocarbons  up  to  butane  are  gases.  From 
pentane  to  hexadecane  they  are  liquids.  The  remainder  are  solids. 
In  composition  each  is  related  to  the  preceding  one  by  containing  the 
additional  units  CH2.  The  formula  of  any  member  of  the  series  is 
therefore  representable  by  the  expression  CnH2n+2.  Substances  re- 
lated in  this  way  form  an  homologous  series.  Their  relations  will  be 
more  clearly  perceived  if  we  employ  the  graphic  formulae.  Since 
hydrogen  appears  uniformly  to  be  univalent,  the  carbon  must  form 
the  backbone  of  each  of  the  compounds.  The  formulae  of  the  first 
members  are  therefore  as  follows  : 


H 

I 

H-C-H 

I 

H 


H    H 

I       I 

H-C-C-H 

I      I 

H    H 


H    H   H 

I       I      I 

H-C-C-C-H 

I       I      I 
H    H    H 


Transferences  of  H  one  step  to  the  right  and  interpositions  of  CH2 
constitute  the  successive  differences. 

A  large  number  of  these  substances  occur  in  nature.  Methane  is 
present  in  large,  and  ethane  in  small,  proportion  in  the  natural  gas  of 
Pennsylvania  and  Ohio.  Many  of  the  others  occur  in  petroleum. 

Petroleum.  —  This  oil  consists  of  a  mixture  of  the  liquid  and 
solid  members  of  the  series  in  varying  proportions,  and  is  found  in 
many  parts  of  the  United  States,  in  Ontario,  at  Baku  on  the  Caspian, 
in  India,  and  in  Japan.  In  oil-refining,  advantage  is  taken  of  the 
differences  in  the  boiling-points  to  make  a  partial  separation  of  the 
components  by  fractional  distillation  (see  below).  The  compounds 
containing  sulphur  which  are  often  present,  and  would  give  the  ob- 
noxious sulphur  dioxide  when  the  oil  was  burned,  are  deprived  of 
this  constituent  by  heating  the  oil  with  powdered  cupric  oxide.  The 
unsaturated  hydrocarbons  (q.v.)  are  removed  by  agitation  with  con- 
centrated sulphuric  acid.  The  following  are  some  of  the  products  of 
the  oil  refinery,  with  their  components  and  uses. 


492 


INORGANIC   CHEMISTRY 


Name. 

Components. 

B.-P. 

Uses 

Petroleum  ether 
Gasolene 
Naphtha 
Benzine 
Kerosene 

Pentane-hexane 
Hexane-heptane 
Heptane-octane 
Octane-nonane 
Decane-hexadecane 

40°-  70° 
70°-  90° 
80°-120° 
120°-150° 
150°-300° 

Solvent,  gas-making 

u                  u 

"         fuel 
n 

Illuminating-oil 

The  portions  of  still  higher  boiling-point  are  employed  as  lubricating 
oils. 

The  vapor  of  these  products  is  more  inflammable  the  more  volatile 
the  components.  The  sale  of  kerosene  is  controlled  legally  by  the  re- 
quirement that  the  vapor  it  gives  when  heated  shall  not  catch  fire 
from  a  naked  flame  until  the  oil  has  reached  a  certain  minimum  tem- 
perature, the  "  flash  point."  This  varies  from  37.7°  to  68.5°  in  different 
states  and  countries. 

By  cooling  the  residues  from  the  retorts  with  a  freezing  mixture 
(cf.  p.  164),  some  of  the  solid  members  of  the  series,  C^H^  to  C28H68, 
are  obtained  as  white  flakes,  which  are  separated  by  filtration  in 
presses.  This  material  forms  the  paraffin  used  in  waterproofing  paper, 
in  laundry  work,  and  as  an  ingredient  in  candles.  In  some  cases 
vaseline,  consisting  of  substances  melting  at  40°-50°,  C^H^  to  CjgH^, 
is  obtained  also. 

From  ozocerite,  which  is  a  sort  of  natural  paraffin,  ceresin,  a  sub- 
stitute for  beeswax,  is  made.  Asphalt  is  another  natural  mixture  of 
hydrocarbons. 

The  formation  of  these  hydrocarbons  in  nature  is  not  yet  thor- 
oughly explained.  According  to  one  theory,  they  are  formed  by  the 
action  of  water  upon  carbides  of  metals ;  while  according  to  another, 
they  result  from  the  decomposition  of  vegetable  or  animal  matter. 
Possibly  both  of  these  sources  have  contributed  to  their  formation. 
Certain  differences  between  the  natural  oils  of  different  localities  point, 
at  all  events,  to  some  difference  in  their  origin. 

Fractional  Distillation.  —  When  the  boiling-points  of  two  com- 
ponents of  a  liquid  are  very  far  apart,  the  vapor  pressure  of  the  one 
may  be  very  low  when  that  of  the  other,  by  heating,  has  reached  760 
mm.  In  this  case  the  first  distillate  will  contain  little  of  the  high- 
boiling  component.  When,  as  in  the  case  of  petroleum,  the  differ- 
ences in  boiling-points  are  not  great,  complete  separation  of  the 


SOME  CARBON  COMPOUNDS  493 

components  is  impossible.  Yet  by  distillation  in  which  the  distillate 
is  caught,  not  in  one  vessel,  but  in  several  successively,  "  fractions  " 
are  obtained  such  that  the  earlier  ones  contain  more  of  the  low-boiling 
and  the  later  ones  more  of  the  high-boiling  materials.  The  vessels 
are  changed  when  the  thermometer  immersed  in  the  vapor  (Fig.  16, 
p.  38)  reaches  certain  temperatures.  When  these  fractions  are  then 
distilled  one  at  a  time,  beginning  with  the  lowest,  and  the  several  dis- 
tillates are  divided  from  one  another  by  the  same  temperatures  as 
before,  a  more  complete  separation  is  effected.  This  process  is  called 
fractional  distillation,  and  may  be  repeated  as  often  as  we  please  with 
constantly  increasing  differentiation  of  the  fractions. 

An  experimental  illustration  may  be  given  by  mixing  0.4  c.c.  of 
benzene  (b.-p.  80.4°)  with  8  c.c.  formic  acid  (b.-p.  100°)  and  2  c.c. 
benzyl  alcohol  (b.-p.  206.5°)  and  boiling  a  part  of  the  mixture  in  a 
test-tube  with  a  small  flame.  The  components  come  off  in  succession, 
and  are  recognized  by  the  fact  that  the  first  and  last  burn  with  a 
luminous  flame,  while  the  flame  of  the  second  is  non-luminous.  By 
passing  the  vapors  into  a  condenser,  and  using  the  method  described 
above,  a  more  or  less  complete  separation  can  be  made. 

General  Properties  of  Hydrocarbons.  —  All  these  substances 
are  extremely  indifferent  in  their  chemical  behavior.  They  have 
none  of  the  properties  of  acids,  bases,  or  salts.  The  halogens,  notably 
chlorine  and  bromine,  however,  interact  with  them  (see  below).  When 
burned  they  all  produce  carbon  dioxide  and  water.  When  their 
vapors  are  passed  through  a  white-hot  tube  they  suffer  decomposition 
into  a  mixture  of  hydrogen  and  hydrocarbons  of  smaller  or  larger 
(see  Benzene)  molecular  weight. 

Methane.  —  Methane,  otherwise  known  as  marsh-gas,  is  the 
chief  component  of  natural  gas.  It  rises  to  the  surface  when  the 
bottoms  of  marshy  pools  are  disturbed,  and  issues  from  seams  in  coal 
beds.  In  these  two  cases  it  results  from  the  decomposition  of  vege- 
table matter  in  absence  of  air.  When  methane  enters  mines  from 
a  coal  seam  it  is  called  "  fire-damp "  (Ger.  Dampf,  vapor),  on  ac- 
count of  the  explosive  nature  of  the  mixture  it  forms  with  the  air. 
The  carbon  dioxide  formed  by  the  explosion  is  called  by  the  miners 
"  choke-damp." 

Methane  may  be  made  from  inorganic  materials  by  the  action  of 


494  INORGANIC   CHEMISTRY 

water  upon  aluminium  carbide,  prepared  by  the  interaction  of  alumi- 
nium oxide  and  carbon  in  the  electric  furnace  (cf.  p.  478) : 

A14C3  +  12H20  ->  4A1(OH)3  +  3CH4. 

In  the  laboratory  the  gas  is  commonly  obtained  by  the  distillation 
of  a  dry  mixture  of  sodium  acetate  and  sodium  hydroxide : 

NaC2H3O2  +  NaOH  ->  Na^COg  +  CH4. 

When  a  mixture  of  methane  and  chlorine  is  exposed  to  sunlight 
several  changes  occur  in  succession  (cf.  pp.  176,  214)  : 

CH4  -f  CL,  ->  CH3C1  +  HC1, 
CH3C1  +  C12  ->  CH2C12  +  HC1, 
O-tLjOljj  -p  Olj  — ^  O-H.O13  -f-  jrLOlj 
CHC18  +  CL,-*  CC14  +HC1. 

This  kind  of  interaction  with  the  halogens  is  characteristic  of  com- 
pounds of  hydrogen  and  carbon.  It  takes  place  slowly,  and  is  there- 
fore entirely  different  from  ionic  chemical  change.  It  consists  in  a 
progressive  substitution  of  chlorine  for  hydrogen,  unit  by  unit.  The 
various  groups  which,  in  the  first  three  of  these  products,  are  asso- 
ciated with  chlorine,  occur  in  many  organic  compounds,  and  receive 
the  names  methyl  (CH8— ),  methylene  (CH2=),  and  methenyl  (CH  =  ). 
The  compounds  are  known,  therefore,  as  methyl  chloride,  methylene 
chloride,  methenyl  chloride  (chloroform),  and  carbon  tetrachloride. 
The  last  two  are  volatile  liquids,  chloroform  being  the  only  one  of  the 
four  which  is  a  familiar  substance.  The  corresponding  iodine  deriva- 
tive iodoform  CHI3  is  a  common  antiseptic.  These  substances  are  not 
salts,  and  are  not  ionized  in  solution.  They  are  very  slowly  hydro- 
lyzed  by  water, — carbon  tetrachloride,  for  example,  giving  carbonic  acid 
and  hydrochloric  acid.  Although  carbon  is  a  non-metal  (cf.  p.  405), 
this  action  requires  a  high  temperature. 

Organic  Radicals.  —  In  carbon  chemistry  there  are  groups  of 
units  which  pass  unaltered  from  compound  to  compound  and  receive 
the  name  organic  radicals.  They  usually  lack  a  property  which  inor- 
ganic radicals  generally  possess,  namely,  the  power  to  form  ions  (p.  312). 
Methyl  is  such  a  radical,  being  found  in  methane  CH3.H,  methyl  chloride 
CH3.C1,  methyl  alcohol  CH3.OH,  and  acetic  acid  CH3.C02H.  Similarly 
we  have  ethyl  C2H6  in  ethane  C2H5.H  and  in  ethyl  alcohol  C2H6.OH. 


SOME  CARBON  COMPOUNDS  495 

Methyl,  ethyl,  and  propyl  (C8H7  —  )  are  univalent  radicals.  We  have 
also  ethylene  C2H4  =  ,  propylene  C8H6  =  ,  and  so  forth,  which  are  biva- 
lent. Groups  like  NO/  (p.  441),  NEL,1  (p.  487),  CHjCO1,  and  many 
more,  are  other  non-ionized  radicals  found  in  organic  compounds  (see 
Acetic  acid,  below). 

Ethylene.  —  In  addition  to  the  paraffin  series  there  are  several 
other  homologous  series  (p.  491)  of  hydrocarbons.  Ethylene  C2H4  is 
the  first  member  of  the  second  series.  It  corresponds  to  ethane,  but 
contains  in  each  molecule  two  hydrogen  units  less  than  does  this 
substance.  The  general  formula  for  this  series  is  C^^.  As  we  shall 
see,  ethylene  and  the  members  of  the  ethylene  series  are  thus  all  un- 
saturated,  possessing  two  free  valences. 

Ethylene  is  most  easily  made  by  heating  common  alcohol,  which 
is  ethyl  alcohol,  with  concentrated  sulphuric  acid  : 

C2H5OH  -*  H20  +  C2H4. 

The  action  really  takes  place  in  two  distinct  stages,  and  the  interme- 
diate product  can  be  isolated.  First,  ethyl  sulphate  (cf.  p.  441)  is 
formed,  C2H6OH  +  H2SO4  <=»  C2H5HSO4  +  H20.  Above  150°,  how- 
ever, this  substance,  which  is  a  thick  syrup,  is  dissociated,  giving 
ethylene  and  sulphuric  acid,  C2H5HSO4  —  »  C2H4  -f  H2S04.  A  compari- 
son of  the  structural  formulse  of  the  alcohol  and  ethylene  shows  that 
this  loss  of  water  must  leave  the  carbon  partly  unsaturated  : 

H    H  H   H  H    H 

II  II  II 

H-C-C-O-H  H-C-C-H     or      H-C  =  C-H 

II  II 

H    H 

The  water  may  also  be  removed  by  allowing  alcohol  to  fall  drop  by 
drop  on  heated  phosphoric  anhydride.  The  solid  phosphoric  acid  re- 
mains behind  and  ethylene  escapes. 

Ethylene  is  formed  along  with  acetylene  and  other  substances, 
when  any  saturated  hydrocarbon  is  heated  strongly.  Even  methane 
gives  it: 

2CH4  -»  C2H4 


Ethylene  is  a  gas,  which,  when  liquefied,  boils  at  —105°.  Its  critical 
temperature  is  130°.  At  0°  it  may  be  liquefied  by  a  pressure  of  42 
atmospheres.  It  burns  in  the  air  with  a  flame  which,  on  account  of 


496  INORGANIC  CHEMISTRY 

the  great  separation  of  free  carbon  which  takes  place  temporarily  dur- 
ing the  combustion  (cf.  Flame),  is  highly  luminous.  It  will  be  seen 
that  in  the  formula  but  three  of  the  valences  of  each  carbon  unit  are 
occupied.  As  carbon  is  either  bivalent  or  quadrivalent,  we  should 
expect  that  in  this  compound  the  combining  capacity  of  the  carbon 
would  not  be  completely  satisfied.  We  find  this  to  be  the  case.  Ethy- 
lene  is  easily  reduced  by  nascent  hydrogen  (p.  423)  to  ethane,  taking 
up  two  units  of  hydrogen  in  the  process.  When  ethylene  is  passed 
through  liquid  bromine  it  is  rapidly  absorbed,  and  the  bromine  seems  to 
increase  in  volume  and  finally  loses  all  its  color,  leaving  a  transparent 
liquid  having  the  composition  C2H4Br2,  ethylene  bromide.  The  second 
of  the  above  graphic  formulae  for  ethylene  is  the  one  generally  used. 
In  spite  of  appearances,  it  is  not  intended  to  indicate  that  the  two 
units  of  carbon  are  more  forcibly  held  together  than  in  other  com- 
pounds (cf.  p.  106).  It  simply  chronicles  the  fact  that  one  valence 
of  each  carbon  unit  is  unoccupied. 

Acetylene.  —  This  substance,  likewise  a  gas,  is  the  first  member  of 
still  another  unsaturated,  homologous  series.  Its  formula  C2H2  shows 
that  its  molecule  lacks  four  of  the  hydrogen  units  necessary  to  the  com- 
plete saturation  which  we  find  in  ethane.  Graphically  its  structure  is 
usually  represented  thus  :  H  —  C  =  C  —  H.  This  gas  is  formed  in 
small  quantities  by  direct  union  of  carbon  and  hydrogen  in  the  electric 
arc  (p.  478).  It  is  also  produced  when  ethylene  is  passed  through 
a  heated  tube  :  C2H4  — >C2H2  +  H2  (cf.  Flame).  When  calcium  carbide 
(p.  478)  is  thrown  into  water,  violent  effervescence  occurs,  the  calcium 
carbide  is  disintegrated,  a  precipitate  of  calcium  hydroxide  is  formed, 
and  acetylene  passes  off  as  a  gas  : 

CaC2  +  211,0  -*  Ca(OH)2  +  C2H2. 

This  action  is  like  that  of  water  on  calcium  phosphide  (p.  461),  calcium 
sulphide  (p.  376),  and  magnesium  nitride  (p.  417). 

Acetylene  burns  with  a  flame  which  is  still  more  luminous  than 
that  of  ethylene.  Its  most  characteristic  property  is  that  when  passed 
through  an  ammoniacal  solution  of  a  cuprous  salt,  it  yields  a  red  pre- 
cipitate of  a  carbide  of  copper  known  as  copper  acetylene.  The 
equation  :  Cu2(OH)2  -f-  C2H2  — >  CugC,,  +  2H20,  partially  represents  the 
change.  This  red  precipitate,  when  dried,  is  extremely  explosive,  en 
account  of  the  great  amount  of  energy  set  free  when  it  breaks  up  into 


SOME   CARBON   COMPOUNDS  497 

its  constituents.     Its  formation  is  used  as  a  test  for  acetylene  in  mix- 
tures of  gases. 

Acetylene  may  be  handled  safely  as  a  gas  at  the  ordinary  pressure, 
but  when  contained  in  cylinders  at  more  than  two  atmospheres  pressure 
it  is  readily  exploded  by  any  shock.  This  is  due  to  the  fact  that  it  is  an 
endothermal  compound : 

CjjHa  -»  2C  +  It,  +  53,200  calories. 

When  used  as  an  illuminant,  it  is  developed  in  a  suitable  generator  as 
it  is  needed.     It  begins  visibly  to  decompose  at  780°. 

The  unsaturated  nature  of  this  substance  is  shown  by  the  avidity 
with  which  it  unites  with  hydrogen  and  the  halogens,  forming  satu- 
rated compounds. 

Benzene.  —  Limits  of  space  forbid  the  discussion  of  any  of  the  other 
series  of  hydrocarbons.  One  of  the  most  important  has  not  been  men- 
tioned, however.  It  is  that  of  which  the  first  member  is  benzene,  C6H6. 
More  than  half  of  the  known  compounds  of  carbon  are  derived  from 
this  substance.  Phenol  (cf.  p.  441)  C6H5OH  is  the  fundamental  alcohol 
of  this  set.  Benzene  is  obtained  from  the  products  of  the  dry  distilla- 
tion of  coal  (cf.  Coal  gas),  being  formed,  probably,  from  the  acetylene 
which  the  decomposition  of  other  hydrocarbons  yields.  At  all  events, 
when  acetylene  is  passed  through  ,a  heated  tube  some  benzene  is  pro- 
duced, 3C2H2  — >  C6H6,  along  with  free  carbon  and  hydrogen. 

THE  ACIDS  OF  CARBON. 

Formic  Acid.  —  The  removal  of  water  from  formic  acid  produces 
carbon  monoxide  (p.  485).  Although  we  cannot  reverse  the  process 
and  cause  carbon  monoxide  to  combine  with  water,  we  can  make  it 
unite  with  bases.  By  passing  carbon  monoxide  over  hot  sodium 
hydroxide,  we  obtain  sodium  formate,  from  which  formic  acid  may  be 
liberated  by  double  decomposition  with  another  acid  : 

CO  +  NaOH  ->  CHO(ONa). 

This  acid  is  secreted  by  red  ants,  and  is  found  in  stinging  nettles.  It 
is  a  liquid  boiling  at  100.1°  and  freezing  at  8.6°.  Although  one  of  the 
weaker  acids,  it  is  much  more  active  than  acetic  acid.  The  molecule 
contains  two  atoms  of  hydrogen,  but  the  acid  is,  in  fact,  monobasic. 


498  INORGANIC   CHEMISTRY 

The  structural  formula  of  the  acid  must  take  account  of  this  fact. 
Three  possibilities  present  themselves  : 

0-H  0  H  0 

C(  II  )C'| 

X0-H  H-C-O-H  H'     ^  0 

In  the  first  and  last  the  hydrogen  units  should  behave  alike.  The 
second  formula  is  the  only  one  which  expresses  the  replaceability  of 
one  unit  and  not  of  the  other  by  a  metal.  Since  the  hydrogen  in 
methane  is  not  replaceable  by  metals  (p.  493),  we  infer  that  the  unit 
directly  combined  with  carbon  is  the  non-replaceable  one.  Sodium 
formate  is  therefore 

0 

II 
H-  C  -  0-Na. 

Acetic  Acid.  —  This  acid  is  produced  in  the  dry  distillation  of 
wood*  (p.  476).  Large  quantities  of  it  are  manufactured  from  dilute 
alcohol.  The  liquid  is  allowed  to  flow  in  a  slow  stream  through  a 
barrel  filled  with  shavings.  Holes  in  the  barrel  provide  for  the  access 
of  air,  and  a  bacterium  with  which  the  shavings  are  infected  promotes 
that  oxidation  of  the  alcohol  in  which  the  change  essentially  consists : 

C2H5OH  -f  02  -*  C2H80(OH)  +  H20. 

Oxygen  alone  does  not  affect  alcohol  in  the  cold.  The  bacterium  (B. 
aceti,  "  mother-of-vinegar  "  )  assists  this  action,  as  lower  organisms  are 
found  to  assist  many  chemical  actions,  in  a  way  which  is  not  as  yet 
thoroughly  understood,  and  which  may  be  described  roughly  as 
catalytic  (see  Fermentation,  p.  501). 

The  dilute  solution  of  acetic  acid  produced  in  this  manner  contains 
from  five  to  thirteen  per  cent  of  acetic  acid,  and  is  known  as  vinegar. 
By  fractional  distillation  the  solution  may  be  concentrated  until  a  little 
water  only  remains,  and  finally,  by  freezing  (cf.  p.  294),  the  acetic  acid 
may  be  crystallized  out.  Pure  acetic  acid,  in  consequence  of  its 
freezing  readily  in  cold  weather,  is  known  as  "  glacial "  'acetic  acid.  It 
melts  at  16.7q  and  boils  at  118°.  . 

Although  four  atoms  of  hydrogen  are  contained  in  its  molecule, 
but  one  of  these  is  replaceable  by  metals.  This  fact  is  recognized  in 
the  constitutional  formula  (p.  391)  of  the  acid,  CH3CO(OH).  In  this 

*  The  dry  distillation  of  bones  (p.  455),  on  the  other  hand,  and  of  animal  mat- 
ter (p.  417)  in  general,  gives  alkaline  liquids,  because  of  the  ammonia  that  is  formed. 


SOME  CARBON  COMPOUNDS  499 

acid  a  radical,  methyl  (CH3— ),  takes  the  place  of  H  in  formic  acid. 
The  other  organic  acids  similarly  are  related  to  formic  acid  and  con- 
tain  other   organic   radicals    (see   Palmitic    acid).     Thus   the   group 
0 

II 
—  C  — O  — H,  called  carboxyl,  is  contained  in  most  carbon  acids,  and  in 

each  of  them,  as  in  formic  acid,  bears  the  replaceable  hydrogen  unit. 
The  other  three  hydrogen  units  in  acetic  acid,  however,  are  replaceable 
by  chlorine,  as  is  the  case  with  the  hydrogen  units  in  hydrocarbons. 

The  above  brief  statements  in  regard  to  the  mode  of  expressing  the  chemical 
properties  of  a  substance  by  an  elaborated  formula  bring  out  a  tendency  which 
prevails  in  the  behavior  of  organic  substances  and  is  almost  entirely  lacking  in 
inorganic  chemistry.  The  units  may  be  removed  from  the  molecule  of  an  organic 
substance  one  by  one,  and  other  units  or  groups  may  be  substituted  for  them  with- 
out disturbing  the  rest  of  the  molecule.  The  changes  take  place,  not  as  in  the  case 
of  ionized  substances,  by  the  splitting  of  the  molecule  into  two  or  more  groups  which 
act  as  wholes,  but  by  the  displacement  of  the  units  piecemeal  and  the  introduction 
of  new  properties  according  to  the  nature  of  the  groups  introduced.  Thus,  if  by  any 

O 
II 

means  we  replace  an  atom  of  hydrogen  by  the  organic  radical— C  —  O  —  H  the  prod- 
uct is  an  acid.  If  we  replace  it  simply  by  the  group  OH  the  product  is  an  alcohol. 
Each  substitution  may  take  place  repeatedly  in  a  given  molecule,  so  that  di-basic 
or  tri-basic  acids,  di-hydric  or  tri-hydiic  alcohols  (see  Glycerine),  or  substances 
which  contain  both  OH  and  — COOH  in  the  same  molecule^like  lactic  acid  and  tar- 
taric  acid),  are  formed.  Other  groups  which  may  be  introduced  or  removed  are 
— NH2,  —  N02,  —  CN,  etc.,  each  of  which  confers  upon  a  substance  the  properties 
which  go  with  the  group,  irrespective  of  the  other  features  which  the  structure  of 
the  substance  may  already  present. 

Oxalic  Acid.  —  This  acid  has  the  composition  H2C204,  and  is  di- 
basic. Its  calcium  salt  is  the  least  soluble  of  the  salts  of  calcium,  and 
is  found  in  many  plants  in  the  form  of  bundles  of  needle-shaped  crys- 
tals. Potassium  hydrogen  oxalate  is  found  in  the  juices  of  various 
species  of  oxalis.  The  acid  may  be  made  by  oxidation  of  sugar  with 
nitric  acid. 

Oxalic  acid  is  commonly  used  in  the  form  of  the  white  crystalline 
hydrate,  H2C2O4, 2H20.  When  heated  carefully  it  sublimes  unchanged. 
Stronger  heating  decomposes  it  into  carbon  dioxide  and  formic  acid, 
and  the  latter  breaks  up,  in  part,  into  water  and  carbon  monoxide.  In 
the  presence  of  dehydrating  agents  like  sulphuric  acid,  water  and  the 
two  oxides  of  carbon  alone  are  formed  (p.  485). 


500  INORGANIC  CHEMISTRY 

CARBOHYDRATES  AND  FERMENTATION. 

Carbohydrates.  —  The  various  kinds  of  sugar,  starch,  and  cellu 
lose  form  a  closely  related  group  of  substances.  As  it  happens  that 
the  proportion  of  hydrogen  to  oxygen  in  the  composition  of  most  of 
them  is  the  same  as  that  of  these  elements  in  water,  they  are  known 
by  the  name  of  carbohydrates.  None  of  these  substances  show  any 
distinct  evidence  of  ionization. 

Dextrose,  otherwise  known  as  glucose  or  grape-sugar,  is  a  white 
crystalline  substance  having  the  composition  C6H1206.  It  is  found 
dissolved  in  the  juices  of  sweet  fruits,  such  as  grapes. 

The  most  familiar  sugar  is  cane-sugar  C12H22On,  which  may  form 
as  much  as  18  per  cent  by  weight  of  the  juices  pressed  from  the 
sugar-cane,  and  sometimes  reaches  15  per  cent  of  the  fluid  material  in 
the  sugar-beet.  It  is  prepared  by  boiling  the  juices  with  animal  char- 
coal (p.  476),  to  remove  the  coloring  matter  which  would  otherwise 
give  the  sugar  a  brown  tint.  The  liquid  is  then  concentrated  until 
crystals  appear.  The  mother-liquor  which  no  longer  deposits  crystals 
is  known  as  molasses. 

When  a  solution  of  cane-sugar  is  boiled  with  water  containing  a 
small  amount  of  any  acid,  a  slow  hydrolysis  of  the  sugar  takes  place, 
whereby  two  other  sugars,  namely,  dextrose  and  levulose,  are  produced  : 
C12H22011  +  H20-»  C6H1206  -f  C6H1206.  The  action  of  the  acid  is  cata- 
lytic, and  the  rate  of  this  more  or  less  leisurely  chemical  change 
depends  upon  the  concentration  of  the  hydrogen  ions.  It  therefore 
furnishes  one  means  of  comparing  acids  as  regards  their  chemical 
activity,  and  has  the  special  advantage  that  the  acid  is  not  consumed 
during  the  process  (of.  p.  356),  but  remains  of  constant  concentration 
throughout  the  whole  time.  This  process,  by  which  cane-sugar  is 
decomposed,  is  spoken  of  as  inversion,  and  the  mixed  product  is  called 
invert-sugar.  The  change  can  be  produced,  not  alone  by  acids,  but 
also  by  certain  complex  chemical  compounds  secreted  by  yeast  (see 
below). 

The  relation  of  starch  to  the  sugars  is  seen,  not  only  in  the  formula 
(CgHjoOg)^  but  in  the  fact  that  by  boiling  starch  with  dilute  acids, 
dextrose  is  formed  along  with  other  products  of  the  hydrolysis.  Com- 
mercial " glucose"  is  made  by  this  process.  Starch  is  an  insoluble 
white  substance  which  is  found  in  the  form  of  fine  particles  in  the 
fruit  and  other  parts  of  plants. 

Cellulose  has  the  same  composition  as  starch.     It  forms  the  frame- 


SOME  CARBON  COMPOUNDS  501 

work  of  the  cells  of  plants.  In  many  cases  it  is  overlaid  with  a  con- 
siderable thickness  of  lignin,  which  in  paper-making  is  removed  by 
boiling  the  wood  with  sodium  hydroxide  or  calcium  bisulphite  solu- 
tion (p.  476).  When  the  product  has  been  washed  thoroughly  with 
water,  almost  pure  cellulose  remains.  Matted  cellulose  in  thin  sheets 
forms  the  basis  of  paper,  and  filter  paper  contains  nothing  else  (see 
under  Aluminium  sulphate).  Other  forms  of  pure  cellulose  are  known 
as  cotton,  linen,  and  jute,  according  to  their  sources.  Although  the 
chemical  composition  of  these  varieties  is  identical,  the  physical  prop- 
erties vary  considerably. 

Fermentation.  —  This  is  the  name  given  to  a  number  of  different 
chemical  changes,  brought  about  by  catalytic  action  of  complex  chemi- 
cal compounds  secreted  by  living  organisms.  These  compounds  are 
called  enzymes,  and,  in  many  cases,  have  been  separated  from  the 
organisms  by  means  of  solvents.  Their  action  must  be  regarded  as 
catalytic,  since  small  quantities  of  the  active  organisms  or  of  the 
enzymes  can  produce  very  extensive  chemical  changes  without  them- 
selves suffering  alteration  in  the  process. 

The  organisms  may  be  divided  into  three  classes,  each  secreting 
different  enzymes  which  confine  themselves  for  the  most  part  to  special 
kinds  of  chemical  change.  (1)  The  molds,  when  grown  in  sugar  solu- 
tion or  beef  extract,  or  other  nutritive  solutions,  produce  decompositions 
known  collectively  as  putrefaction.  (2)  Certain  bacteria  promote  the 
oxidation  of  alcohol  to  acetic  acid  (p.  498).  Some  also  decompose 
sugar,  furnishing  butyric  or  lactic  acid  as  one  of  the  products.  (3) 
The  yeasts  (saccharomycetes)  flourish  in  solutions  of  some  sugars,  and 
decompose  them  into  alcohol  and  carbon  dioxide.  This  decomposition 
is  known  as  alcoholic  fermentation.  These  changes  are  usually  brought 
about  by  actual  introduction  of  the  organism.  In  brewing,  however, 
the  enzyme  itself,  diastase,  is  employed  to  hydrolyze  starch. 

The  juice  of  grapes  when  set  aside  at  a  suitable  temperature  soon 
begins  to  ferment,  owing  to  the  propagation  in  it  of  a  yeast  (S.  ellipsoi- 
deus)  which  is  found  upon  the  skins  of  the  grapes,  and  decomposes  the 
sugar.  While  small  quantities  of  a  number  of  different  compounds  are 
formed,  by  far  the  greater  part  of  the  sugar  is  resolved  quantitatively 
into  alcohol  and  carbon  dioxide  :  C6H1206  — >  2C2H6OH  +  2CO2.  The 
liquid  effervesces,  and  the  carbon  dioxide  escapes  into  the  air.  The 
wine  is  allowed  to  stand,  after  fermentation,  until  it  has  deposited  a 
considerable  crust  of  material  known  as  argol,  which  consists  mainly  of 


502  INORGANIC  CHEMISTRY 

potassium-hydrogen  tartrate  (KHC4H4O6,  cream  of  tartar).  The  con- 
centration of  the  sugar  in  the  grape-juice  being  small,  the  quantity  of 
alcohol  contained  in  the  product  is  not  very  great.  By  distillation  of 
wine,  a  liquid  containing  a  much  larger  proportion  of  alcohol  is  made, 
and  is  known  as  brandy.  The  special  flavors  of  wines  and  brandies 
depend  upon  materials,  other  than  sugar,  originally  contained  in  the 
fermented  liquid,  upon  by-products  of  the  fermentation,  and  upon 
materials  which  arise  by  slow  chemical  changes  while  the  liquor  is 
stored. 

The  preparation  of  beer  involves  a  preliminary  step,  for  the  sugar 
needed  for  the  fermentation  is  made  from  the  starch  contained  in  various 
kinds  of  grain,  particularly  barley.  The  conversion  of  starch  into 
sugar  is  effected  by  the  use  of  diastase,  which  is  formed  in  considerable 
amounts  in  sprouting  barley.*  When  the  quantity  of  this  unorganized 
ferment  has  reached  a  maximum,  the  barley  is  dried  and  crushed.  The 
product,  known  as  malt,  is  mixed  in  water  with  other  grain  which  has 
not  gone  through  this  process,  and  the  whole  is  slowly  heated.  During 
the  heating,  the  diastase  hydrolyzes  the  molecules  of  the  starch,  giving 
a  sugar  called  maltose  (C^H^On).  The  solution,  after  filtration  and 
boiling,  is  cooled.  It  is  then  placed  in  the  fermentation  vats,  and  com- 
mon yeast  ($.  cerevisice)  is  added.  During  the  growth  of  this  plant  the 
maltose  is  decomposed,  each  molecule  producing  two  molecules  of 
dextrose,  and  the  latter  is  broken  up  into  alcohol  and  carbon  dioxide 
according  to  the  equation  already  given.  Aside  from  the  alcohol  and 
carbon  dioxide,  considerable  quantities  of  other  substances  extracted 
from  the  grain  remain  in  the  solution  and  form  the  so-called  "  extract," 
which  varies  in  kind  and  quantity  in  different  varieties  of  beer. 

Whiskey  is  prepared  in  a  somewhat  similar  fashion,  although  other 
sources  of  starch,  such  as  rye  and  corn,  may  be  used.  The  fermented 
liquid  contains  but  a  small  proportion  of  alcohol,  and  is  distilled 
(rectified).  The  alcohol,  being  more  volatile  than  water,  tends  to  pass 
over  first,  and  a  product  containing  any  desired  proportion  of  it  (see, 
however,  below)  can  be  made.  Alcohol  for  technical  or  chemical  use  is 
made  in  the  same  way,  potatoes  sometimes  being  used  as  a  source  of 
the  starch. 

*  In  the  digestion  of  bread,  potatoes,  and  other  food  containing  starch,  the 
same  office  is  performed  by  the  saliva.  Starch,  being  insoluble,  could  not  be 
absorbed  through  the  walls  of  the  alimentary  tract,  while  sugar  is  soluble  and  can  be 
so  absorbed.  In  a  roughly  similar  way  the  albuminous  parts  of  food  are  rendered 
soluble  and  capable  of  assimilation  by  the  pepsin  of  the  stomach,  and  the  pan- 
creatin  of  the  large  intestine,  which  convert  them  into  peptones. 


SOME   CARBON   COMPOUNDS  503 

ALCOHOLS,  ESTERS,  AND  ETHERS. 

Alcohols.  — We  have  already  seen  that  when  wood  is  distilled, 
methyl  alcohol  is  found  in  the  fluid  product.  When  purified  this  is  a 
colorless  liquid  boiling  at  66°.  Its  solution  in  water  shows  no  evidence 
of  ionization,  although  it  would  probably  be  safer  to  say  that  the 
ionization  is  so  slight  as  to  be  imperceptible,  than  to  say  that  the  com- 
pound is  not  ionized  at  all.  The  formula  (CH8OH)  makes  it  impossible 
to  represent  the  structure  of  the  substance  in  more  than  one  way  : 

H 

I 
H-C-O-H 

I 
H 

All  alcohols  contain  the  group  =C— O— H  (cf.  p.  499). 

Common  alcohol,  ethyl  alcohol  C2H5OH,  is  formed  in  the  fermenta- 
tion of  solutions  of  sugar  by  yeast  (p.  501),  and  is  separated  from  the 
water  and  the  other  products  of  fermentation  by  distillation.  The 
product  contains  95  per  cent  of  alcohol  and  5  per  cent  of  water,  and 
is  applicable  to  most  commercial  uses.  Absolute  alcohol,  entirely  free 
from  water,  cannot  be  made  by  distillation  alone  (see  below).  The 
95  per  cent  spirit  is  placed  in  vessels  filled  with  quicklime,  the  latter 
interacts  with  water  producing  calcium  hydroxide,  and  the  clear  liquid 
which  is  poured  off  is  distilled  once  more.  Pure  alcohol  boils  at  78.3°. 

Mixtures  of  two  liquids,  when  distilled,  behave  in  one  of  three  ways. 
Two  of  these  have  been  described  already  (p.  183),  and  alcohol  (b.-p. 
78.30°)  and  water  (b.-p.  100°)  illustrate  the  third.  In  this  case  the 
vapor  tension  of  a  certain  mixture  is  higher  than  that  of  any  other 
mixture  and  higher  than  that  of  either  component  separately.  This 
special  mixture  has,  therefore,  a  lower  boiling-point  than  any  other. 
In  the  present  instance  this  mixture  contains  95.57  per  cent  of  alcohol 
and  4.43  per  cent  of  water  and  boils  at  78.15°.  When  the  fermented 
liquid,  with  its  large  percentage  of  water,  is  distilled,  the  alcohol  all 
tends  to  pass  off  first,  in  association  with  that  part  of  the  water  required 
to  constitute  the  mixture  of  minimum  boiling-point.  Eepeated  distil- 
lation simply  eliminates  more  completely  the  excess  of  water  beyond 
this  amount  (viz.,  4.43  per  cent),  by  leaving  it  in  the  residues. 

Glycerine  (p.  452)  is  an  alcohol  containing  three  hydroxyl  groups, 
a  trihydric  alcohol. 


504  INORGANIC  CHEMISTRY 

Esters.  —  When  an  organic  acid  and  an  alcohol  are  mixed,  a  very 
slow  chemical  action  takes  place,  which,  being  reversible,  in  no  case 
reaches  completion.  With  the  simplest  members  of  these  groups,  for- 
mic acid  and  methyl  alcohol,  for  example,  the  change  is : 

HCOOH  +  HOCH3  <=»  HCOOCH8  +  ILf). 

The  product  is  known  as  methyl  formate.  The  corresponding  action 
between  acetic  acid  and  ethyl  alcohol :  CH3COOH  +  HOC2H6  <±  CH3CO 
OC2H5  -f-  H20,  results  in  the  formation  of  ethyl  acetate.  In  this  case, 
when  equivalent  quantities  of  the  initial  substances  have  been  used 
without  any  solvent,  and  a  condition  of  equilibrium  has  been  reached, 
two-thirds  of  the  material  is  found  to  have  been  transformed  into 
ethyl  acetate  and  water.  If  we  start  with  the  latter  materials  in  pure 
form,  the  same  equilibrium  point  is  reached,  and  one-third  of  the  mate- 
rial is  converted  into  acetic  acid  and  alcohol. 

This  action  is  a  general  one,  and  occurs  between  all  alcohols  and 
acids.  The  products  are  sometimes  known  as  ethereal  salts,  because 
they  result  from  the  displacement  of  the  hydrogen  of  an  acid  by  a  radi- 
cal. This  designation,  however,  is  not  very  happy,  since  the  products 
are  not  ionized  and  possess  none  of  the  properties  of  salts.  The 
special  name  eaters,  therefore,  has  been  given  to  them.  The  action  is 
always  extremely  slow  and  never  complete,  but  it  may  be  hastened  and 
carried  to  completion  by  the  introduction  of  some  substance  capable  of 
absorbing  the  water  and  so  preventing  the  reversal.  Concentrated  sul- 
phuric acid,  for  example,  or  anhydrous  cupric  sulphate,  may  be  used. 

Inorganic  acids  also  interact  with  alcohol,  giving  esters.  Thus, 
nitroglycerine  (p.  452)  is  an  ester,  and  should  be  called  glyceryl  trini- 
trate.  The  use  of  sulphuric  acid  to  assist  in  the  removal  of  the  water 
is  illustrated  in  the  preparation  of  this  substance.  Gun-cotton  (p.  441) 
is  an  ester  of  nitric  acid  also,  for  cellulose  is  a  complex  alcohol.  Ethyl 
sulphate  (p.  495)  is  an  ester  of  sulphuric  acid.  In  this  case  the  action 
may  be  made  complete  by  using  sulphuric  acid  containing  an  amount 
of  sulphur  trioxide  sufficient  to  combine  with  the  water  to  be  pro- 
duced. 

The  above  actions,  in  which  an  ester,  like  ethyl  acetate,  is  formed, 
may  be  almost  completely  reversed  if  a  sufficient  amount  of  water  is 
added  (cf.  p.  250).  The  hydrolysis  of  the  ester  is  hastened  by  the 
presence  of  free  acids  in  the  water.  This  is  owing  to  the  catalytic 
action  of  the  hydrogen  ions,  and  the  acceleration  is  proportional  to  the 
activity  of  the  acid  used.  The  acid,  however,  although  it  hastens  the 


SOME  CARBON  COMPOUNDS  505 

action,  does  not  carry  it  beyond  the  condition  of  equilibrium  which  it 
would  eventually  have  reached  with  the  same  amount  of  water  alone 
(cf.  p.  237). 

When  esters  are  boiled  with  strong  bases,  such  as  sodium  hydroxide 
solution,  the  salt  of  the  acid  and  an  alcohol  are  formed  : 

CH8COOC2H5  +  NaOH  -^  CH8COONa  -{-  HOC2H6. 

With  more  complex  esters  the  sodium  salts  of  the  acids  thus  produced 
are  known  as  soaps,  and  this  general  kind  of  action  is  called,  therefore, 
saponification  (Lat.  sapo,  soap).  The  speed  with  which  it  proceeds 
may  be  used  as  a  means  of  measuring  the  activity  of  bases. 

Soap.  —  Soap  is  prepared  by  the  decomposition  of  fat.  The  latter 
substance  is  a  mixture  of  several  rather  complex  esters.  In  beef  fat 
the  chief  esters  present  are  tripalmitin,  tristearin,  and  triolein.*  It 
will  be  sufficient  to  illustrate  the  chemistry  of  the  change  by  discussing 
the  case  of  one  of  these  substances.  Tripalmitin  is  the  glyceryl  ester  of 
palmitic  acid.  When  fat  is  mixed  with  hot  sodium  hydroxide  solution 
it  first  forms  an  emulsion  f  in  which  the  fat  is  disseminated  in  minute 
droplets  through  the  liquid.  This  is  a  result  of  surface  tension. 
When  the  emulsion  is  boiled,  the  fat  is  slowly  decomposed  into  sodium 
palmitate  and  glycerine.  The  change  is  precisely  similar  in  plan  to 
the  simpler  one  just  discussed  : 

C15H31COO  -  C  -  H2  HOCH2 

I  I 

C15H31COO  -  C  -  H     H-  3NaOH  -»  SC^COONa  +  HOCH 

I  I 

C15H31COO  -  C  -  H2  HOCH2 

Changes  similar  to  this  occur  with  the  other  two  substances.  The 
only  difference  is  that  the  organic  radical  in  the  case  of  tristearin  is 
C^HgB,  and  in  the  case  of  triolein  C^Hgg.  Both  products  in  each  case  are 
soluble  in  water,  but  when  common  salt  is  added  to  the  solution  the 
sodium  salts  of  the  organic  acids  are  separated  ("salted  out,"  see 
Chap,  xxxv)  as  a  solid  mass,  which  is  known  as  soap.  When  potas- 
sium hydroxide  takes  the  place  of  sodium  -hydroxide  the  mass  is  semi- 
fluid, and  is  known  as  soft  soap. 

*  Butter  fat  contains  in  addition  to  the  above  a  certain  amount  of  tributyrin, 
in  which  the  organic  radical  is  C3H7.  Olive  oil  consists  mainly  of  tripalmitin  and 
triolein. 

t  In  the  intestines  the  same  office  is  performed  by  the  gall,  secreted  by  the  liver, 
and  so  the  fat  is  prepared  for  absorption  into  the  system. 


506  INORGANIC   CHEMISTRY 

These  complex  esters,  like  the  simpler  ones,  are  decomposed  also 
by  water  without  the  aid  of  a  base,  although  much  more  slowly.  By 
the  use  of  superheated  steam,  however,  rapid  hydrolysis  can  be  pro- 
duced, and  the  products  are  the  free  organic  acids  and  glycerine.  The 
mixture  of  acids  set  free  from  fat  by  the  action  of  steam  is  a  solid, 
waxy  mass,  known  as  "  stearin,"  and  used  in  the  manufacture  of 
candles.  The  oleic  acid,  which  is  a  liquid,  is  pressed  out. 

These  acids,  not  being  soluble  in  water,  have  no  effect  upon  litmus  ; 
but  the  fact  that  they  are  acids* may  be  recognized  when  it  is  found 
that  they  are  converted  into  soluble  salts  by  bases,  such  as  sodium 
hydroxide : 

O^H^COOH  I  H-  NaOH  <=>  H20  +  C17H85COONa. 

The  cleansing  power  of  soap  solution  seems  to  depend  on  the  surface 
tension  of  the  liquid  rather  than  on  any  chemical  action. 

The  components  of  soap,  like  other  salts,  are  highly  ionized  in  solu- 
tion, and  show  all  the  properties  of  ionogens.  When  soap  is  dissolved 
in  hard  water  (cf.  p.  113),  a  white,  flocculent  precipitate  is  formed,which 
coagulates  upon  the  sides  of  the  vessel.  This  is  a  mixture  of  the  cal- 
cium salts  formed  by  union  of  the  proper  ions.  For  example,  the 
sodium  palmitate  is  changed  as  follows  : 

2C15H31COONa  +  CaS04  ->  (C16H81COO),Ca{+  Na2S04. 

Most  of  the  salts  of  these  acids,  with  the  exception  of  those  of  potas- 
sium and  sodium,  are  insoluble  in  water. 

Drying  Oils,  —  The  oils  commonly  used  as  "  dryers  "  for  mixing 
with  varnish  and  paint  and  in  making  linoleum,  such  as  linseed  oil, 
hemp  oil,  poppy  oil,  and  nut  oil,  contain  esters  of  acids  with  unsaturated 
radicals.  One  of  the  constituents,  for  example,  is  the  glyceryl  ester 
of  linoleic  acid.  The  formula  of  this  acid  is  C17H31COOH.  It  con* 
tains  four  hydrogen  atoms  less  than  the  corresponding  saturated  acid 
(stearic  acid).  These  oils,  especially  after  having  been  recently 
heated,  alone  or  with  catalytic  agents  like  lead  oxide  and  manganese 
dioxide,  absorb  oxygen  rapidly  from  the  air,  and  become  solid.  They 
do  not  dry,  in  the  ordinary  sense,  by  evaporation. 

Ether.  —  When  two  molecules  of  an  alcohol  lose  one  molecule  of 
water,  an  ether  is  produced : 

2CH3OH  ->  (CH3)20  +  H20. 


SOME  CARBON  COMPOUNDS  507 

Thus,  methyl  alcohol  gives  methyl  ether,  and  ethyl  alcohol,  ethyl  or 
common  ether.  The  action  is  most  easily  carried  out  by  two  steps. 
In  making  common  ether,  ethyl  alcohol  acts  upon  sulphuric  acid,  giv- 
ing ethyl  sulphate  (p.  495)  ;  and  the  latter,  when  warmed  gently  with 
excess  of  alcohol,  gives  ethyl  ether  : 

C2H6HS04  +  C2H5OH  liii  (C2H6)20 1  +  H2S04. 

The  ether  escapes  as  vapor  and  is  condensed. 

Ethyl  ether  is  a  volatile  liquid  boiling  at  34.6°.  It  is  largely  used 
as  a  solvent  for  iodine,  fats,  and  other  substances  not  readily  soluble 
in  water,  and  as  an  anaesthetic. 

CYANOGEN. 

Cyanogen.  —  This  compound  is  formed  in  small  amount  when  a 
discharge  of  electricity  takes  place  between  carbon  poles  in  an  atmos- 
phere of  nitrogen  (cf.  p.  416).  Cyanogen,  being  an  endothermal  sub- 
stance, is  more  easily  made  as  one  product  in  an  exothermal  action 
(p.  301).  It  is  prepared  by  allowing  a  solution  of  cupric  sulphate  to 
trickle  into  a  warm  solution  of  potassium  cyanide.  The  cupric  cyanide, 
at  first  precipitated,  quickly  decomposes,  giving  cuprous  cyanide  and 
cyanogen : 

2KNC  +  CuS04  -»  Cu(NC)2J  +  I^SO^ 

2Cu(NC)2  ->  2CuNC  +  C2N2  f . 

Cyanogen  is  a  very  poisonous  gas  with  a  characteristic,  faint  odor. 

Hydrocyanic  Acid.  —  This  acid,  called  also  prussic  acid,  has  the 
formula  H  —  N  =  C,  and  is  most  easily  made  by  the  action  of  an  acid 
upon  a  cyanide  (see  Potassium  cyanide)  followed  by  distillation.  It 
is  a  colorless  liquid  boiling  at  26.5°.  It  has  an  odor  like  that  of  bitter 
almonds,  and  is  highly  poisonous.  In  aqueous  solution  it  is  an  ex- 
tremely feeble  acid,  and  is  hardly  ionized  at  all.  In  consequence  of 
this,  potassium  cyanide  is  markedly  hydrolyzed  by  water,  and  its  aque- 
ous solution  is  strongly  alkaline.  The  behavior  of  hydrocyanic  acid 
shows  it  to  be  an  unsaturated  body,  a  fact  which  is  taken  account  of 
in  the  above  formula,  and  illustrated  in  the  two  following  paragraphs. 

Cyanates.  —  When  potassium  cyanide  is  fused  and  stirred  with 
an  easily  reducible  oxide,  like  lead  oxide  (PbO),  the  metal  (for  example, 


508  INORGANIC  CHEMISTRY 

the  lead)  collects  at  the  bottom  of  the  iron  crucible  in  molten  form,  and 
potassium  cyanate  is  produced: 

KNC  +  PbO  ->  KNCO  +  Pb. 

Cyanic  acid  is  itself  very  unstable.  Ammonium  cyanate  is  chiefly 
remarkable  for  its  transformation  into  urea  (p.  488). 

Thiocyanates.  —  When  potassium  cyanide  in  aqueous  solution  is 
boiled  with  sulphur  or  with  a  polysulphide  (p.  376),  it  is  converted  into 
potassium  thiocyanate,  KCNS.  This  salt  is  used  in  testing  for  ferric 
ions  on  account  of  the  deep-red  color  of  ferric  thiocyanate  (cf.  p.  250). 
The  ammonium  salt  undergoes  at  170°  a  transformation  parallel  to 
that  of  ammonium  cyanate,  thiocarbamide  (sulpho-urea)  being  formed. 

Exercises.  —  1.  Make  the  graphic  formulae  of  hexane  (p.  491), 
methyl  acetate  (p.  499),  ethyl  formate,  ethylene  bromide  (p.  496), 
oxalic  acid  (p.  499),  ethyl  ether  (p.  507). 

2.  Make    equations    for    the    formation    of    aluminium    carbide 
(p.  478),  the  hydrolysis  of  starch  to  maltose   (p.  502),  the  saponifi- 
cation   of  triolein   (p.  505). 

3.  Name  the  radicals  :  C6Hn,  C6H10,  C5H9,   C16H13,   and  the  sub- 
stances C6HUC1,   C6HnOH,   (C6Hn)20,   C5HUHSO4. 

4.  Is  the   hydrocarbon  radical   of  oleic   acid  (p.  505)  saturated, 
or  not  ? 

5.  Prepare  a  summary  of  the  various  statements  that  have  been 
made  in  the  text  about  catalysis    (e.g.  pp.  75,  109,  111,  170,  175,  189, 
237,  380,   396,  399,  423,  484,  500,  501),  and  illustrate  fully. 


CHAPTER   XXX 
FLAME 

Meaning  of  the  Term.  —  In  the  combustion  of  charcoal  there  is 
hardly  any  flame,  for  the  light  emanates  almost  entirely  from  the  incan- 
descent, massive  solid.  When  two  gases  are  mixed  and  set  on  fire,  a  sort 
of  flame  passes  through  the  mixture,  but  this  can  hardly  be  accounted  a 
flame,  in  the  ordinary  sense,  either.  The  rapid  movement  of  the  flash, 
and  the  explosion  which  accompanies  it,  are  in  a  manner  the  precise 
opposite  of  the  quiet  combustion  which  is  characteristic  of  flames. 

With  illuminating-gas  the  production  of  its  very  characteristic  flame 
is  due  to  the  chemical  union  of  a  stream  of  one  kind  of  gas  in  an 
atmosphere  of  another.  The  flame  is  made  up  of  the  heated  matter 
where  the  two  gases  meet.  In  the  case  of  a  burning  candle,  one  of  the 
bodies  appears  to  be  a  solid,  but  a  closer  scrutiny  of  the  phenomenon 
shows  that  the  solid  does  not  burn  directly.  A  combustible  gas 
is  manufactured  continuously  by  the  heat  of  the  combustion  and 
rises  from  the  wick.  The  introduction  of  a  narrow  tube  into  the 
interior  of  the  flame  enables  us  to  draw  off  a  stream  of  this  gas  and  to 
ignite  it  at  a  remote  point.  Thus,  a  flame  is  a  phenomenon  produced 
at  the  surface  where  two  gases  meet  and  undergo  combination  with 
the  evolution  of  heat  and,  sometimes,  light. 

In  the  chemical  point  of  view,  it  is  a  matter  of  indifference  whether 
the  gas  outside  the  flame  contains  oxygen,  and  the  gas  inside  consists 
of  substances  ordinarily  known  as  combustibles,  or  whether  this  order 
is  reversed.  In  an  atmosphere  of  ordinary  illuminating-gas,  the  flame 
must  be  fed  with  air.  This  condition  is  easily  realized  (Fig.  93). 
The  lamp-chimney  is  closed  at  the  top  until  it  has  become  filled  with 
illuminating-gas.  After  the  lapse  of  a  few  minutes  this  can  be  ignited 
as  it  issues  from  the  bottom  of  the  wide,  straight  tube  which  projects 
from  the  interior.  When  the  hole  in  the  cover  of  the  lamp-chimney  is 
then  opened,  the  upward  draft  causes  the  flame  of  the  burning  gas  to 
recede  up  the  tube,  and  there  results  a  flame  fed  by  air  and  burning  in 
coal-gas.  In  an  atmosphere  of  this  kind,  materials  playing  the  part  of 
a  candle  burning  in  air  would  have  to  be  substances  which,  under  the 

509 


510 


INORGANIC  CHEMISTRY 


influence  of  the  heat  of  combustion,  give  off  oxygen.  Strongly  heated 
potassium  chlorate,  for  example,  appears  to  burn  continuously  in  such 
an  atmosphere  when  lowered  into  it  in  a  deflagrating  spoon. 

Luminous  Flames.  —  The  flame  of  hydrogen,  under  ordinary  cir- 
cumstances, is  almost  invisible,  nearly  all  the  energy  of  the  combustion 
being  devoted  to  the  production  of  heat.  A  part  of  this,  however, 
may  be  transformed  into  light  by  the  sus- 
pension of  a  suitable  solid  body,  such  as  a 
platinum  wire,  in  the  flame.  The  holding  of 
a  piece  of  quicklime  in  an  oxy hydrogen  flame 
(cf.  p.  109)  is  a  practical  illustration  of  this 
method  of  securing  luminosity.  In  general, 
luminosity  may  be  produced  by  the  presence 
of  some  incandescent  solid. 

In  the  Welsbach  lamp  the  flame  itself  is 
non-luminous,  and,  but  for  the  mantle,  would 
be  identical  with  the  ordinary  Bunsen  flame. 
The  mantle  which  hangs  in  the  flame,  how- 
ever, by  its  incandescence,  furnishes  the  light. 
This  mantle  is  composed  of  a  mixture  of  99 
per  cent  thorium  dioxide  (Th02)  and  one  per 
cent  cerium  dioxide  (Ce02).  While  many 
oxides  would  give  out  a  white  light  and  could 
be  obtained  much  more  cheaply  than  these, 
they  have  not  sufficient  coherence  to  make 
their  use  practicable.  It  is  worth  noting  that 
any  appreciable  variation  from  the  above  pro- 
portions, by  the  introduction  of  either  more 
or  less  cerium  oxide,  produces  a  marked 
diminution  in  the  intensity  and  whiteness  of 
the  light  (see,  also,  under  Thorium). 

In  cases  of  brilliant  combustion,  as  of  magnesium  ribbon  or  phos- 
phorus, a  solid  body  is  formed  whose  incandescence  accounts  for  the 
light.  The  flame  of  ordinary  illuminating-gas  does  not  at  first  sight 
appear  to  give  evidence  of  the  presence  of  any  solid  body.  But  if  a 
cold  evaporating  dish  is  held  in  the  flame  for  a  moment,  a  thick  deposit 
of  finely  divided  carbon  (soot)  is  formed,  and  we  at  once  realize  that 
the  light  is  due  to  the  glow  of  these  particles  in  a  mass  of  intensely 
hot  gas.  Carbon  is,  indeed,  an  extremely  combustible  substance,  and 


FIG.  93. 


FLAME  511 

is  eventually  entirely  consumed.  But  a  fresh  supply  is  continually 
being  generated  in  the  interior  of  the  flame,  while  the  oxygen  with 
which  it  is  to  unite  is  outside  the  flame  altogether.  Thus  the  carbon 
particles  persist  until,  drifting  with  the  spreading  gas,  they  reach  the 
periphery  of  the  flame. 

It  cannot  be  said  that  no  flames  are  luminous  unless  a  solid  body  is 
contained  in  them.  When  compressed  hydrogen  is  burned  in  an  atmos- 
phere of  oxygen  under  pressure,  the  light  given  out  by  the  flame  is 
much  greater,  and,  in  general,  illuminating  power  seems  to  be  height- 
ened by  increase  in  the  density  of  the  gas.  In  special  cases,  also,  such 
as  the  explosion  of  a  mixture  of  nitric  oxide  with  carbon  bisulphide 
vapor,  a  flame  which  has  considerable  illuminating  power  is  produced, 
although  the  density  of  the  gases  is  low  and  solids  are  lacking. 

The  'JBunsen  Flame.  —  In  the  burner  devised  by  Robert  Bunsen,  a 
jet  of  ordinary  illuminating-gas  is  projected  from  a  narrow  opening 
into  a  wider  tube.  In  this  tube  it  becomes  mixed  with  air,  drawn  in 
through  openings  whose  dimensions  can  be  altered  by  means  of  a  per- 
forated ring.  When  the  supply  of  air  is  sufficient,  the  flame  becomes 
non-luminous.  With  a  somewhat  different  construction,  and  the  use  of 
a  bellows  to  force  a  larger  proportion  of  air  into  the  gas,  a  still  hotter 
flame  can  be  produced.  The  instrument  in  this  case  is  known  as  a 
blast-lamp.  The  high  temperature  of  the  Bunsen  flame  is  not  difficult 
to  account  for.  It  will  be  seen  at  once,  on  handling  the  burner,  that 
the  flame  diminishes  to  one-half  or  one-third  of  its  previous  size  when 
air  is  admitted.  Since  the  same  amount  of  gas  is  burning  in  both 
cases,  and  the  products  in  both  cases  are  ultimately  alike,  the  total 
amounts  of  heat  produced  must  in  both  cases  be  the  same.  This  state- 
ment may  safely  be  made,  since  experiment  shows  that  when  chemical 
substances  pass  from  one  condition  to  another,  the  amount  of  heat 
evolved  is  the  same  whatever  intermediate  steps  may  be  taken  or 
omitted  (cf.  p.  78).  The  production  of  an  equal  amount  of  heat  in  a 
smaller  flame  necessarily  causes  this  to  have  a  higher  average  temper- 
ature. It  will  be  noted  that  it  does  not  follow  from  this  that  the  tem- 
perature is  higher  in  every  part  of  the  non-luminous  flame  than  in  the 
corresponding  locality  in  the  luminous  flame.  We  shall  see  later,  in- 
deed, that  this  is  not  the  case. 

It  is  instructive  to  note  the  effect  of  forcing  in  larger  and  larger 
proportions  of  air  into  the  Bunsen  flame.  The  flame  at  the  top  of  the 
tube  continually  diminishes  in  size,  even  after  it  has  become  non- 


512 


INORGANIC  CHEMISTRY 


luminous.  Finally,  a  point  is  reached  at  which  the  flame  is  so  unstable 
that  the  smallest  further  increase  in  the  supply  of  air  causes  it  to 
descend  into  the  tube.  The  mixture  of  illuminating-gas  and  air  in  the 
tube  of  a  Bunsen  burner  is  an  explosive  one,  and  the  explosion-flame 
will  proceed  through  it  with  greater  rapidity,  the  more  nearly  the 
quantity  of  air  approaches  that  required  for  complete  combustion. 
When  the  speed  with  which  the  explosion-flame  would  move,  equals 
that  with  which  the  stream  of  the  mixed  gases  is  coming  upward 
through  the  tube,  the  flame  reaches  the  unstable  condition  just  men- 
tioned. Any  increase  in  the  proportion  of  air  raises  the  speed  with 
which  the  explosion  can  travel,  and  the  flame  is  thus  able  to  proceed 
down  the  tube  against  the  stream  of  gas.  This  phenomenon  is  fre- 
quently noticed  in  the  Bunsen  burner,  when  the  holes  admitting  the 
air  are  too  large,  or  a  draft  momentarily  causes  an  increase  in  the 
supply  of  air.  The  flame  "  strikes  back,"  and  thereafter  continues 
to  burn  at  the  bottom  of  the  tube. 


Structure  of  the  Bunsen  Flame.  —  When  an  ex- 
ceedingly small  luminous  flame  is  examined,  the  various 
parts  of  which  it  consists  may  easily  be  made  out.  In  the 
interior  there  is  a  dark  cone  which  is  composed  of  illumi- 
nating-gas and  air,  and  in  it  no  combustion  is  taking  place. 
A  match-head  may  be  held  here  for  some  time  without 
being  set  on  fire.  This  is  therefore  not  properly  a  part  of 
the  flame.  Outside  this  is  a  vivid  blue  layer  ((7,  Fig.  94) 
which  is  best  seen  in  the  lower  part  of  the  flame,  but  ex- 
tends beneath  the  luminous  sheath,  and  covers  the  dark 
inner  cone  completely.  Outside  the  blue  flame,  and  cover- 
ing the  greater  part  of  it,  is  the  cone-shaped  luminous 
portion  (B).  Over  all  is  an  invisible  mantle  of  non-lumi- 
nous flame  (A),  which  becomes  visible  when  the  light  from 
the  luminous  part  is  purposely  obstructed.  In  the  lumi- 
nous gas-flame,  therefore,  there  are  four  regions,  if  we 
count  the  inner  cone  of  gas.  The  difference  between  this 
and  the  non-luminous  Bunsen  flame  is  that  in  the  latter 
the  luminous  region  is  omitted,  and  the  inner,  dark  cone,  the  blue 
sheath,  and  the  outer  mantle,  are  the  only  parts  which  can  be  dis- 
tinguished. We  shall  see  that  in  these  different  regions  the  chemical 
changes  taking  place  are  different. 


FIG.  94. 


FLAME 


518 


Composition  of  Illuminating -Gas.  —  Before  considering  the 
chemistry  of  the  gas-flame,  it  is  necessary  to  know  what  substances  are 
burning.  The  illuminating-gas  in  Europe,  and  in  many  of  the  smaller 
cities  of  the  United  States,  is  usually  coal-gas,  while  in  the  larger  cities 
of  America  it  is  almost  always  made  from  water-gas.  The  former  is 
obtained  by  the  destructive  distillation  of  soft  coal,  and  is  freed  from 
ammonia  (cf.  p.  418)  and  tar  by  washing  and 'cooling,  and  from  hydrogen 
sulphide  and  carbon  dioxide  by  passage  through  layers  of  slaked  lime. 
The  water-gas,  made  by  the  action  of  steam  upon  anthracite  or  coke, 
being  composed  of  carbon  monoxide  and  hydrogen  (cf.  p.  485),  has  no 
illuminating  power.  It  is  therefore  "  carburetted,"  that  is,  mixed  with 
hydrocarbons,  by  passage  through  a  cylindrical  structure  filled  with 
white-hot  firebrick,  upon  which  falls  a  small  stream  of  high-boiling 
petroleum.  The  relatively  involatile  hydrocarbons  of  which  the  oil  is 
composed  are  thus  decomposed  ("cracked"),  and  gaseous  substances 
of  high  illuminating  power  are  produced.  The  following  table  shows 
the  composition  of  each  of  these  kinds  of  gas,  together  with  that  of 
oil-gas  (Pintsch's),  which  is  composed  entirely  of  the  products  from 
"  cracking  "  oil  : 


Components. 

Coal-Gas. 

Water-Gas. 

Oil-Gas. 

Illuminants  .... 

6.0 

16.6 

45.0 

Heating  gases: 

Methane  .... 

34.5 

19.8 

38.8 

Hydrogen 

49.0 

32.1 

14.6 

Carbon  monoxide     . 

7.2 

26.1 

Impurities: 

Nitrogen  .... 

3.2 

2.4 

1.1 

Carbon  dioxide   .     . 

1.1 

3.0 

.  .  . 

Candle  power    . 

17.5 

26.0 

65.0 

These  are  average  numbers,  and  considerable  variations  from  these  pro- 
portions are  often  met  with.  The  illuminants  are  unsaturated  hydro- 
carbons, such  as  ethylene,  acetylene,  and  benzene,  and  the  value  of  the 
gas  for  illuminating  purposes  depends  on  the  amount  of  these  particular 
components. 


Chemical    Changes   Taking  Place  in  the  Bunsen  Flame.— 

The  study  of  the  chemical  changes  taking  place  in  the   Bunsen  flame, 


514  INORGANIC   CHEMISTRY 

particularly  with  the  object  of  explaining  (1)  the  luminosity  of 
the  flame  of  the  pure  gas,  and  (2)  the  non-luminosity  of  that 
produced  by  the  same  gas  when  it  is  mixed  with  air,  has  been  the  sub- 
ject of  many  elaborate  investigations.  The  questions  are  :  Why  is 
carbon  liberated  in  the  former  case,  and  why  is  it  not  liberated  in  the 
latter  ?  Let  us  consider  these  questions  in  order. 

1.  The  first  suggestion  of  an  explanation  of  the  presence  of  free 
carbon  in  the  luminous  flame  was  that  the  hydrocarbons  contained  in 
the  gas  underwent  a  selective  combustion.  This  was  supposed  to  take 
place  in  such  a  way  that  the  hydrogen  was  first  burned  out  of  the 
molecules  by  the  oxygen,  which,  by  diffusion,  had  penetrated  farthest 
into  the  flame.  The  carbon  had,  therefore,  to  remain  free  until  all  the 
hydrogen  was  satisfied,  and  until  the  former  had  drifted  into  a  region 
in  which  a  more  liberal  supply  of  oxygen  was  obtainable.  Thus  the 
carbon  was  rendered  incandescent  by  the  heat  of  combustion  of  the 
hydrogen.  This  hastily  formed  theory  is  founded  upon  ill-considered 
premises,  and  is  not  supported  by  experiment.  It  is  based  upon  the 
familiar  fact  that  hydrogen  gas  burns  more  easily  than  solid  charcoal. 
Such  a  comparison,  however,  is  not  applicable  to  the  case  of  the  com- 
bustion of  a  gaseous  hydrocarbon,  for  in  such  a  substance  both  of  the 
elements  are  in  a  gaseous  condition,  and  the  inferior  combustibility  of 
charcoal  is  due,  not  to  an  inferior  power  of  uniting  with  oxygen,  but 
to  the  fact  that,  since  it  is  a  solid,  the  oxygen  finds  access  to  it  only 
with  difficulty. 

The  question  could  manifestly  be  placed  beyond  dispute  if  it  were 
found  possible  to  determine  exactly  the  nature  of  the  change  which 
passage  through  the  inner,  blue  zone  of  the  flame  produced.  If,  for 
example,  gases  could  be  extracted  by  means  of  a  tube  from  the  region 
just  outside  this  zone,  and  it  were  found  that  all  the  hydrogen  had  been 
converted  into  water  while  the  carbon  had  not  yet  reached  the  stage  of 
carbon  dioxide,  this  theory  would  be  confirmed.  It  was  found  impos- 
sible, however,  to  perform  the  experiment  in  this  fashion,  on  account 
of  the  disturbing  influence  of  the  currents  produced  by  the  withdrawal 
of  the  stream  of  gas.  An  apparatus  devised  by  Smithells,  however, 
furnished  an  ingenious  means  of  securing  the  interconal  gas,  alone  and 
in  any  quantity.  The  air  and  combustible  gas  are  admitted  in  propor- 
tions which  can  be  varied,  and  the  mixture  burns  at  the  top  of  the 
wider  tube  (Fig.  95).  As  the  quantity  of  air  is  increased,  however,  the 
speed  with  which  an  explosion-flame  would  pass  through  it  becomes 
greater,  and  finally  the  inner  cone  passes  down  and  rests  upon  the 


FLAME 


515 


mouth  of  the  narrow  tube  through  which  the  mixture  of  gases  is  issuing 
more  rapidly.     A  preliminary  combustion  takes  place  in  the  blue  cone, 
while  the  final  conversion  of  the  whole  material  into  carbon  dioxide  and 
water  is  completed  in  the  outer  mantle.     This 
remains  at  the  top  of  the  wider  tube,  where 
alone  the  necessary  air  can  be  obtained.     By 
means  of  a  side-tube  (not  shown)  fused  into 
the  wider  tube,  the  gases  which  have  traversed 
the  inner  cone  and  undergone  chemical  change 
in  it  can  be  extracted  and  subjected  to  exam- 
ination. 

Now  it  was  found  that  hardly  any  of  the 
hydrogen  had  been  burned  by  the  blue  cone, 
while  all  of  the  carbon  had  reached  the  stage 
of  carbon  monoxide.  This  was  the  case,  not 
only  with  illuminating-gas,  which  initially 
contains  much  free  hydrogen,  but  also  when 
the  flame  was  fed  with  pure  methane.  The 
greater  part  of  the  hydrogen  of  the  latter 
was  set  free,  and  was  found  uncombined  in 
the  interconal  gas.  It  is  evident,  therefore, 
that  when  a  hydrocarbon  undergoes  partial 
combustion  the  carbon  is  first  attacked  and 
the  burning  of  the  hydrogen  is  postponed 
until  a  sufficient  supply  of  oxygen  becomes 
available. 

We  may  safely  infer,  therefore,  that  since 
in  the  luminous  flame  a  portion  at  least  of 
the  carbon  remains  for  a  time  uncombined, 

the  hydrogen  must  a  fortiori  have   survived  combustion  also.     The 
luminous  cone  must  contain  a  mixture  of  free  hydrogen  and  free  carbon. 

The  production  of  this  mixture  can  be  explained  in  only  one  way, 
namely,  by  the  dissociation  of  the  hydrocarbons  concerned,  of  which 
ethylene  is  the  most  important.  Now  we  have  seen  (pp.  496  and  497) 
that,  when  heated,  ethylene  dissociates,  giving  acetylene,  and  acetylene 
gives  carbon  and  hydrogen.  Hence  this  stochastic  hypothesis  is  highly 
probable.  It  only  remains  to  be  shown  that  acetylene  is  actually  formed 
as  an  intermediate  substance.  Evidence  of  this  is  obtained  in  two  ways. 
It  is  found  that  when  the  Bunsen  flame  "  strikes  back,"  and  the  combus- 
tion of  the  gases  is  rendered  incomplete  by  the  contact  of  the  flame  with 


FIG.  95. 


516 


INORGANIC   CHEMISTRY 


the  cold  tube,  a  large  amount  of  acetylene  is  formed.  Again,  when  the 
gases  surrounding  the  flame  of  air  burning  in  illuminating-gas  (p.  509) 
are  withdrawn  by  means  of  a  pump  and  caused  to  pass  through  an 
ammoniacal  solution  of  cuprous  chloride  (Fig.  96),  large  quantities 
of  copper  acetylene  are  formed. 

The  conception  that  when  hydrocarbons  burn,  they  first  undergo  dis- 
sociation, and  then  union  with  oxygen,  is  in  harmony  with  what  we 

have  observed  also  in  the 
case  of  the  combustion  of 
hydrogen  sulphide,  where 
the  presence  of  free  sul- 
phur and  free  hydrogen 
in  the  interior  of  the  flame 
was  demonstrated  (p. 
373). 

2.  The  influence  of 
the  air  admitted  to  the 
Bunsen  burner,  in  inter- 
fering with  this  dissocia- 
tion in  such  a  way  as  to 
destroy  all  luminosity,  is 
FIG.  96.  the  most  difficult  point  to 

explain.      The    effect    is 

frequently  attributed  to  the  oxygen  which  the  air  contains.  This  view, 
however,  is  seriously  weakened  by  a  consideration  of  the  undoubted 
fact  that  oxygen  is  not  required.  Carbon  dioxide  and  steam  are  equally 
efficient  when  introduced  instead  of  air.  Even  nitrogen,  which  cannot 
possibly  be  suspected  of  furnishing  any  oxygen,  likewise  destroys  the 
luminosity.  Lewes  has  shown  that  0.5  volumes  of  oxygen  in  1  volume 
of  coal-gas  destroy  the  luminosity.  But  2.30  volumes  of  nitrogen  or 
2.27  volumes  of  air  accomplish  the  same  result.  Thus  the  efficiency 
of  air  is  not  much  greater  than  that  of  nitrogen,  in  spite  of  the  fact 
that  one-fifth  of  the  former  is  oxygen. 

It  is  evident  that  the  effect  is  due,  in  part  at  least,  to  the  dilution 
with  cold  gas.  This  is  confirmed  by  the  observation  that  a  cold  plati- 
num dish  held  in  a  small  luminous  flame  is  similarly  destructive  of  the 
luminosity.  Comparison  of  the  temperatures  of  the  inner  sheaths  of 
the  luminous  and  non-luminous  flames  shows  that  the  temperature  of 
the  latter  is  markedly  lower.  If  the  tube  of  the  Bunsen  burner  is 
heated  so  that  the  mixed  gases  are  considerably  raised  in  temperature 


FLAME  517 

before  reaching  the  non-luminous  flame,  the  latter  becomes  luminous. 
It  is  probable,  therefore,  that  the  cold  gas  lowers  the  temperature  of  the 
inner  flame,  and  at  the  same  time  the  dilution  diminishes  the  speed  of 
dissociation  (Lewes).  Even  if  the  temperature  is  not  reduced  below 
that  at  which  dissociation  of  the  ethylene  can  occur,  yet  the  dilution 
and  cooling  together  prevent  that  sharp  dissociation  at  this  particular 
point  which  is  necessary  for  the  production  of  the  great  excess  of  free 
carbon  needed  to  furnish  the  light. 

Exercises.  —  1.  In  what  way  will  calcium  hydroxide  remove 
hydrogen  sulphide  from  coal-gas  (p.  513)  ? 

2.  Make  a  section  showing  the  shape  of  the  flame  produced  by 
burning  hydrogen  gas  when  the  latter  issues  from  a  circular  opening. 


CHAPTER   XXXI 
SILICON  AND   BORON 

IN  respect  to  chemical  relations  there  is  a  close  resemblance  between 
silicon  and  carbon.  The  former  element  gives  no  monoxide,  however, 
and  is  quadrivalent  in  all  its  compounds.  In  chemical  character  it  is 
strictly  non-metallic. 

Occurrence.  —  Silicon,  unlike  carbon,  is  not  found  in  the  free  con- 
dition. In  combination  it  is  the  most  plentiful  element  after  oxygen, 
and  constitutes  more  than  one-quarter  of  the  crust  of  the  earth.  The 
oxide  is  silica  or  sand  (Si02),  and  this  oxide  and  its  compounds  are 
components  of  many  rocks.  In  the  inorganic  world  silicon  is  the  char- 
acteristic element  to  almost  as  great  an  extent  as  is  carbon  in  the 
organic  realm. 

Preparation.  —  When  finely  powdered  magnesium  and  sand  are 
mixed,  and  one  part  of  the  mass  is  heated,  a  violent  action  spreads 
rapidly  through  the  whole  : 

2Mg  +  Si02  ->  Si  +  2MgO. 

At  the  same  time,  and  especially  if  excess  of  the  metal  is  used,  some 
magnesium  silicide  Mg2Si  is  formed  also.  The  mixture  is  treated  with 
a  dilute  acid  which  decomposes  the  magnesium  oxide  and  the  silicide, 
and  leaves  the  silicon  alone  undissolved.  A  simpler  method,  although 
one  using  less  common  materials,  is  to  pass  the  vapor  of  silicon  tetra- 
chloride  over  heated  sodium.  The  mixture  of  common  salt  and  silicon 
which  remains  in  the  tube  can  be  separated  by  washing  with  water. 
Both  of  these  methods  give  amorphous  silicon  in  the  state  of  fine  powder. 
When  amorphous  silicon  is  dissolved  in  molten  zinc,  the  mass,  after 
solidification,  is  found  to  contain  crystalline  silicon.  This  form  may 
be  made  in  one  operation,  by  heating  three  parts  of  potassium  fluosili- 
cate  K2SiF6  with  one  part  of  sodium  and  four  parts  of  zinc  in  a  closed 
crucible.  The  sodium  displaces  the  silicon  and  combines  with  the 
fluorine,  while  the  zinc  acts  as  a  solvent  as  before.  The  zinc  is  removed 
lay  the  action  of  a  dilute  acid,  the  silicon  remaining  unaffected. 

518 


SILICON  AND  BORON  519 

Properties.  —  Amorphous  silicon  is  a  brown  powder.  It  unites 
with,  fluorine  at  the  ordinary  temperature,  with  chlorine  at  430°,  with 
bromine  at  500°,  with  oxygen  at  400°,  with  sulphur  at  600°,  with  nitro- 
gen at  about  1000°,  and  with  carbon  and  boron  at  temperatures  attain- 
able only  in  the  electric  furnace.  It  is  slowly  oxidized  by  aqua  regia 
to  silicic  acid,  and  is  dissolved  by  a  mixture  of  hydrofluoric  acid  and 
nitric  acid,  giving  the  fluoride. 

Crystallized  silicon  consists  of  black  needles  belonging  to  the  hex- 
agonal system,  and  is  less  active  than  the  other  variety.  It  oxidizes 
superficially  at  400°,  and  the  dioxide  formed  on  the  surface  hinders 
further  action.  With  chlorine  and  fluorine  it  unites  easily  when 
heated.  Gaseous  hydrogen  fluoride  interacts  violently  with  it  at  a  high 
temperature,  heat  and  light  are  given  out,  and  silicon  tetrafluoride  and 
hydrogen  are  formed.  It  is  slowly  attacked  by  hydrofluoric  acid 
mixed  with  nitric  acid,  but  not  by  any  others  of  the  oxygen  acids. 

Both  kinds  of  silicon  act  with  fair  speed  upon  boiling  solutions  of 
potassium  or  sodium  hydroxide  (cf.  p.  99),  the  metasilicate  being 
formed : 

Si  +  2KOH  +  H20  rfe  K2Si03  +  2H2. 

Silicon  seems  to  be  more  active  than  carbon,  for,  when  it  is  heated  with 
fused  potassium  carbonate,  potassium  silicate  is  formed  and  carbon  is 
liberated.  Both  kinds  of  silicon  differ  from  carbon  in  being  fusible  at 
a  very  high  temperature.  The  product  is  crystallized  silicon. 

Silicon  Hydride.  —  Silicon  differs  from  carbon  in  giving  only  two 
well-defined  compounds  with  hydrogen.  The  chief  one  may  be  liberated 
as  a  gas  by  the  action  of.  hydrochloric  acid  upon  magnesium  silicide : 

Mg2Si  +  4HC1  -»  2MgCLj  +  SiH4. 

The  action  is  similar  to  that  by  which  hydrogen  sulphide  is  made. 
Since  the  magnesium  silicide  always  contains  free  magnesium,  hydro- 
gen is  liberated  at  the  same  time.  The  gases  may  be  separated  by 
passage  through  a  tube  surrounded  by  liquid  air.  The  silicon  hydride 
is  reduced  to  liquid  form,  while  the  hydrogen  passes  on.  The  mixture 
with  hydrogen  is  spontaneously  inflammable.  The  pure  gas  becomes 
so  only  when  its  pressure  is  reduced.  In  the  air,  however,  it  is  easily 
inflammable,  by  contact  with  a  warm  body.  When  heated,  it  decom« 
poses  into  its  constituents. 


520  INORGANIC   CHEMISTRY 

Carbide  of  Silicon.  —  This  compound  is  manufactured  for  use  as 
an  abrasive,  and  is  sold  under  the  name  of  carborundum.  A  mixture 
of  quartz-sand,  coke,  and  common  salt  is  heated  to  about  3500°  in  the 
electric  furnace  : 

Si02  +  30  _>  SiC  +•  2CO. 

The  materials  are  piled  upon  an  oblong  base  provided  with  vertical 
walls  at  the  ends,  through  which  the  terminals  project.  A  long  ridge 
of  loosely  packed  coke  furnishes  an  imperfect  conductor  of  high  resist- 
ance between  the  terminals.  The  above  mixture  is  heaped  on  each  side 
of  and  above  this,  as  well  as  below  it.  The  coke  becomes  white-hot 
when  the  current  is  turned  on,  and  the  heat  radiating  from  this  core 
brings  about  a  change,  in  the  above  sense,  which  affects  the  whole  mass 
more  or  less  completely. 

Silicon  carbide  when  pure  is  composed  of  transparent,  colorless, 
hexagonal  plates.  Ordinarily  the  crystals  are  brown  or  black.  It 
stands  next  to  the  diamond  and  carbide  of  boron  in  hardness.  It  is 
not  oxidized  by  the  air  even  at  a  white  heat,  being  protected  by  the 
layer  of  silica  first  formed.  Carborundum  is  not  affected  by  water  or 
acids,  but  is  decomposed  by  alkalies.  It  is  used  in  making  machinery 
for  polishing  hard  rock,  such  as  granite,  in  making  whetstones  and 
grindstones,  and  is  employed  also  for  protecting  the  walls  of  puddling 
furnaces  (q.v.),  and  in  other  ways  in  the  steel  industry. 

Silicon  Tetrachloride.  —  This  compound  is  made  by  direct  union 
of  the  free  elements.  It  is  more  conveniently  prepared  by  passing 
chlorine  over  a  strongly  heated  mixture  of  silicon  dioxide  and  carbon. 
The  gaseous  products  enter  a  condenser  in  which  the  tetrachloride 
assumes  the  liquid  form  : 

2CI,  +  Si02  H-  20  -»  SiCl4  -f  2CO. 

Chlorine  is  unable  to  displace  oxygen  from  combination  with  silicon, 
and  has,  therefore,  when  alone,  no  effect  upon  sand.  In  the  above 
action,  therefore,  the  carbon  is  used  to  secure  the  oxygen  while  the 
chlorine  combines  with  the  silicon.  This  kind  of  interaction  is  in 
some  degree  different  from  any  which  we  have  hitherto  encountered. 
It  bears  no  special  name,  but  the  principle  underlying  it  is  very  com- 
monly employed  (see,  e.g.,  Chlorides  of  boron  and  aluminium). 

Silicon  tetrachloride  is  a  colorless  liquid  (b.-p.  59°)  which  fumes 


SILICON  AND  BORON  521 

strongly  in  moist  air,  giving  silicic  acid.  It  acts  violently  upon  cold 
water,  and  in  this  respect  differs  from  carbon  tetrachloride : 

SiCl4  +  4H20  ->  4HC1  +  Si(OH)4. 

The  silicic  acid  (q.v.)  which  is  formed  soon  appears  as  a  gelatinous  pre- 
cipitate. 

When  silicon  is  heated  in  a  stream  of  dry  hydrogen  chloride,  a 
mixture  of  silicon  tetrachloride  and  silico-chloroform  SiHClg  is  pro- 
duced. The  latter  is  a  volatile  liquid  boiling  at  34°. 

Silicon  Tetrafluoride.  —  When  strong  hydrofluoric  acid  acts  upon 
sand,  this  gas  is  liberated : 

Si02  +  4HF  ->  2H20  +  SiF4. 

Since  the  water  interacts  with  the  tetrafluoride  (see  below),  the  latter 
is  usually  made  by  the  use  of  powdered  calcium  fluoride  and  excess 
of  sulphuric  acid.  In  this  way  the  hydrogen  fluoride  is  generated  in 
contact  with  the  sand,  and  at  the  same  time  the  sulphuric  acid  takes 
possession  of  the  water.  Hydrofluoric  acid  acts  in  a  corresponding 
way  upon  all  silicates  (q.v.),  whether  these  are  minerals  or  are  artificial 
silicates  like  glass  (cf.  p.  243). 

Silicon  tetrafluoride  is  a  gas  which  becomes  solid,  without  liquefy- 
ing (cf.  p.  463),  when  cooled  to  —102°.  It  fumes  strongly  in  moist 
air,  and  acts  vigorously  upon  water.  This  interaction  is  different  from 
that  of  the  tetrachloride,  because  the  excess  of  the  tetrafluoride  forms 
a  complex  compound  with  the  hydrofluoric  acid  : 

SiF4  +  4H20  -»  Si(OH)4  (  +  4HF)  (1) 

(4HF)  +  2SiF4  ->  2H2SiF6 (2) 

3SiF4  -f  4H2O  ->  Si(OH)4  +  2H2SiF6 

The  silicic  acid  is  precipitated  in  the  water,  and  may  be  separated  by 
filtration,  leaving  a  solution  of  hydrofluosilicic  acid. 

Hydrofluosilieic  Acid.  —  This  acid  is  stable  only  in  solution. 
When  the  water  is  removed  by  evaporation,  silicon  tetrafluoride  is  given 
off,  while  most  of  the  hydrogen  fluoride  remains  to  the  last.  Its  salts 
are  decomposed  in  a  corresponding  way  when  they  are  heated.  This 
acid  is  used  in  analysis  chiefly  because  its  potassium  and  sodium  salts 
are  amongst  the  few  salts  of  these  metals  which  are  relatively 


522  INORGANIC  CHEMISTRY 

insoluble  in  water.  The  barium  salt  is  also  insoluble,  but  most  of  the 
salts  of  the  heavy  metals  are  soluble. 

Silicon  Dioxide.  —  This  substance  may  be  made  in  the  form  of  a 
fine  white  powder  by  heating  precipitated  silicic  acid.  It  is  found  in 
many  different  forms  in  nature.  In  large,  transparent,  six-sided  prisms 
with  pyramidal  ends  it  is  known  as  quartz  or  rock  crystal.  When  col- 
ored by  manganese  and  iron  it  is  called  amethyst,  when  by  organic 
matter,  smoky  quartz.  A  special  arrangement  of  the  structure  gives 
cat's  eye.  Amorphous  forms  of  the  same  material,  often  colored  brown 
or  red  with  ferric  oxide,  are  agate,  jasper,  and  onyx,  the  last  much  used 
in  making  cameos.  Infusorial  or  diatomaceous  earth  is  composed  of  the 
tests  of  minute  organisms.  Slightly  hydrated  forms  of  silica  are  the 
opal  and  flint.  A  crystalline  variety  belonging  to  the  hexagonal  sys- 
tem, but  showing  entirely  different  crystalline  forms,  occurs  occasion- 
ally in  minute  crystals,  and  is  called  tridymite. 

Silicon  dioxide,  although  differing  profoundly  from  carbon  dioxide 
in  its  physical  nature,  nevertheless  behaves  like  the  latter  chemically. 
Thus,  when  boiled  with  potassium  hydroxide  solution  (cf.  p.  482),  it 
forms  potassium  orthosilicate : 

Si02  +  4KOH  ->  K4Si04  +  2H20. 

The  salt  is  left  as  a  gelatinous  solid  ("soluble  glass  ")  when  the  water 
is  evaporated.  The  silicates  of  potassium  and  sodium  may  also  be 
obtained  by  boiling  sand  with  the  carbonates  of  these  metals,  carbon 
dioxide  being  displaced.  They  are  produced  more  rapidly,  however, 
as  metasilicates  (see  below),  by  fusing  the  sand  with  the  alkali 
carbonates  : 

Si02  +  K2C03  i»  K2Si03  +  C02. 

Silica  is  found  in  the  hard  parts  of  straw,  of  some  species  of  horse- 
tail (equisetum),  and  of  bamboo.  In  the  form  of  whetstones  it  is  used 
for  grinding.  The  clear  crystals  are  employed  in  making  spectacles 
and  optical  instruments.  Pure  sand  is  used  in  glass  manufacture 
(g.v.).  Infusorial  earth,  on  account  of  the  tubular  form  of  many  of  the 
minute  structures  of  which  it  is  composed,  can  absorb  three  times  its 
own  weight  of  nitroglycerine,  and  is  therefore  employed  in  making 
dynamite.  Kecently,  small  pieces  of  chemical  apparatus  have  been 
manufactured  by  fusing  quartz  in  the  oxyhydrogen  flame.  The  mate- 
rial is  very  difficult  to  work,  on  account  of  its  inf usibility ;  but  owing 


SILICON  AND  BORON  523 

to  the  low  coefficient  of  expansion  of  silica,  the  vessels  fashioned  out 
of  it  can  be  heated  or  cooled  as  suddenly  as  we  choose,  without  risk  of 
fracture. 

Silicic  Acid*  —  When  acids  are  added  to  a  solution  of  sodium  sili- 
cate, silicic  acid  is  set  free.  After  a  little  delay  it  usually  appears  as  a 
gelatinous  precipitate.  When,  however,  the  silicate  is  poured  into 
strong  hydrochloric  acid,  no  precipitation  occurs.  The  silicic  acid 
remains  in  colloidal  solution  (see  below).  The  acid  before  precipi- 
tation is  supposed  to  be  orthosilicic  acid  : 


4HC1    *  4NaCl 

2HC1  +  H2O  ->  2NaCl  +  Si(OH)4, 

but  the  gelatinous  precipitate  when  it  has  been  dried  contains  a 
smaller  proportion  of  the  elements  of  water.  There  seem  to  be  no 
definite  stages,  indicating  the  existence  of  various  acids,  such  as  we 
observe  with  phosphoric  acid.  We  should  expect  the  vapor  tension  of 
the  water  to  decrease  by  steps,  each  of  which  should  correspond  to  some 
acid  of  a  particular  degree  of  hydration  (cf.  p.  122),  but  nothing  of  the 
sort  is  observed.  The  final  product  of  drying  is  the  dioxide.  A  jelly 
of  the  kind  described  above  is  called  a  hydrogele.  A  solution  such  as 
that  containing  dissolved  silicic  acid  is  called  a  hydrosole. 

Silicic  acid  is  a  very  feeble  acid,  and,  therefore,  gives  no  salt  with 
ammonium  hydroxide.  For  the  same  reason  silicic  acid  can  be  com- 
pletely displaced  from  its  salts,  even  by  so  weak  an  acid  as  carbonic  acid. 
Since  the  water  in  some  geyser  regions  contains  alkali  silicates,  the  action 
of  the  carbon  dioxide  in  the  air  causes  deposition  of  silica  round  the 
places  where  the  water  issues  from  the  ground.  This  form  is  known 
as  silicious  sinter.  Striking  scenic  effects,  as  in  the  white  and  pink 
terraces  of  New  Zealand,  are  sometimes  produced  by  this  method  of 
deposition. 

Colloidal  Solution.  —  Silicic  acid  in  colloidal  solution  can  be 
separated  from  the  other  dissolved  substances  by  a  method  dis- 
covered by  Graham  and  called  dialysis.  The  solution  is  placed  in 
a  vessel  whose  bottom  is  composed  of  vegetable  parchment,  or  some 
animal  membrane,  and  the  whole  is  hung  in  a  vessel  of  water.  Salts 
which  may  be  present  pass  slowly  through  the  membrane  into  the 
water  outside,  and,  when  the  latter  is  frequently  changed,  all  of  the 
material  of  this  class  in  the  solution  may  finally  be  removed. 


524  INORGANIC   CHEMISTRY 

The  silicic  acid,  however,  remains  confined  in  the  cell.  Substances 
possessing  the  ability  to  pass  through  a  membrane  are  usually  crystal- 
line, and  in  the  present  connection  are  therefore  called  crystalloids. 
On  the  other  hand,  materials  of  a  gluey  nature  are  unable  to  traverse 
the  partition,  and  are  named  colloids.  This  behavior  must  not  be 
confused  with  that  in  which  a  semi-permeable  membrane  (p.  284)  is  in 
question.  In  the  latter  case  all  dissolved  substances  are  equally 
unable  to  penetrate  the  membrane. 

A  pure  aqueous  solution  of  a  colloid  lacks  the  properties  charac- 
teristic of  solutions.  It  boils  and  freezes  at  the  same  temperature  as 
the  pure  solvent,  and  independent  evidence  shows  that,  although  in 
too  fine  a  state  of  division  to  be  retained  by  a  filter  paper,  the  colloid 
is  nevertheless  simply  suspended  in  the  liquid. 

Colloidal  solutions  of  other  substances  which  are  insoluble  in  water, 
such  as  ferric  hydroxide,  aluminium  hydroxide,  and  starch,  and  even 
of  metals  like  gold,  silver,  and  platinum,  can  likewise  be  made. 

Silicates.  —  While,  in  the  absence  of  definite  knowledge,  silicic 
acid  is  presumed  to  be  the  ortho-acid  Si(OH)4,  and  no  other  silicic  acids 
have  been  made,  the  salts  are  most  easily  classified  by  imagining  them 
to  be  derived  from  various  acids  representing  different  degrees  of  hy- 
dration  of  the  dioxide,  or,  to  put  it  the  other  way,  different  degrees  of 
dehydration  of  the  ortho-acid.  The  following  equations  show  the  rela- 
tion of  the  ortho-acid  to  some  of  the  silicic  acids  whose  salts  are  most 
commonly  found  amongst  minerals  : 

H4SiO4  -  H20  ->  H2Si03     Metasilicic  acid. 
2H4Si04  -  H20  ->  H6Si207  1  -, 
2H4Si04-3H2O^H2Si205  |DlSlllC1C  aClds' 
3H4Si04-  4H2O  ->  H4Si308    Trisilicic  acid. 

Di-  andtrisilicates  are  those  derived  from  acids  containing  two  and  three 
units  of  silicon,  respectively,  in  the  formula. 

The  composition  of  minerals  is  often  exceedingly  complex.  This  is 
due  to  the  fact  that  amongst  them  mixed  salts  (p.  359)  are  exceedingly 
common,  in  which  the  hydrogen  of  the  imaginary  acid  is  displaced  by 
two  or  more  metals  in  such  a  way  that  the  total  quantity  of  the  metals 
is  equivalent  to  the  hydrogen.  The  following  list  presents  in  tabular 
form  some  typical  or  common  minerals  arranged  according  to  the 
above  mentioned  classification : 


SILICON  AND   BORON  525 

Zircon,     ZrSi04 


Orthosilicates  (H4Si04)  . 


Garnet,  Ca8Fe2rn(Si04)8 
Mica,  KH2Al3(SiO4)8 
Kaolin,  H2AL2(Si04)2,H2O 


f  Wollastonite,     CaSiO; 
Metasilicates  (H2Si08) 


Enstatite,  8 

Hornblende  and  augite 

Disilicate  (H6Si2O7)          Serpentine,  Mg3Si207, 
Trisilicate  (H4Si808)         Orthoclase  (feldspar),  KAlSi3O8 

It  will  be  seen  that  the  total  valence  of  the  metal  units  is  equal  to  that 
of  the  acid  radicals.  Thus,  in  beryl  there  are  six  equivalents  of 
glucinum  (beryllium)  and  six  of  aluminium,  taking  the  place  of  twelve 
units  of  hydrogen  in  (H2Si03)6. 

Mica,  which  is  obtained  in  large  sheets  from  Farther  India,  is  used 
in  making  lamp-chimneys  and  as  an  insulator  in  electrical  apparatus. 
Kaolin,  or  clay,  like  mica,  is  an  acid  orthosilicate.  It  is  formed  in 
nature  as  the  result  of  the  action  of  water  and  carbonic  acid  upon 
minerals  like  feldspar.  In  such  cases,  those  elements  which  can  form 
carbonates,  like  potassium,  magnesium,  and  calcium,  are  usually  dis- 
placed from  combination  with  the  silicic  acid.  Aluminium,  however,  is 
too  feeble  a  base  to  form  a  carbonate,  and  is  thus  left  in  combination 
as  silicate.  A  clay  containing  lime  is  called  a  marl,  and  one  contain- 
ing sand,  a  loam. 

Some  of  these  minerals  frequently  occur  mixed  together  as  regular 
components  of  certain  igneous  rocks.  Thus,  granite  is  a  more  or  less 
coarse  mixture  of  quartz,  mica,  and  feldspar.  Frequently  the  oblong, 
flesh-colored  or  white  crystals  of  the  last  are  large  and  very  conspicu- 
ous both  in  granite  and  in  porphyry.  In  basalt  the  components  are 
usually  less  easily  visible  to  the  eye.  Lava  is  the  name  for  any 
rock  recently  ejected  from  a  volcano.  Pumice-atone  is  the  name 
given  to  the  parts  of  the  lava  which  are  porous,  having  acquired  this 
texture  from  the  expansion  of  bubbles  of  gas  consequent  upon  release 
of  pressure.  Sandstone  is  composed  of  sand  cemented  together  by  clay 
or  by  calcium  carbonate,  and  colored  brown  or  yellow  by  ferric  oxide. 

Since  many  silicates  are  not  affected  by  acids,  or  at  least  are  affected 
with  extreme  slowness,  special  means  have  to  be  taken  to  get  the  con- 
stituents of  such  minerals  into  soluble  form  for  the  purpose  of  analysis, 


526  INORGANIC   CHEMISTRY 

Two  methods  are  in  use.  Sometimes  the  finely  powdered  mineral  is 
heated  in  a  platinum  dish  with  hydrofluoric  acid  until  all  the  silicon 
has  passed  off  in  the  form  of  the  tetrafluoride.  Since  the  use  of  the 
fluorides  would  lead  to  difficulties  in  the  course  of  the  analysis,  the 
resulting  mixture  of  fluorides  of  the  metals  is  next  heated  strongly 
with  concentrated  sulphuric  acid,  and  the  mixture  of  sulphates  thus 
produced  is  treated  according  to  the  usual  routine.  In  other  cases  the 
finely  powdered  mineral  is  fused  at  a  white  heat  with  a  mixture  of 
potassium  and  sodium  carbonates.  In  this  way  carbonates  of  the 
metals  are  formed,  along  with  potassium  and  sodium  silicate.  Treat- 
ment with  water  dissolves  the  latter,  and  leaves  the  carbonates  to  be 
handled  in  the  usual  way. 

BORON. 

As  regards  chemical  relations,  boron,  being  a  uniformly  trivalent 
element,  is  a  member  of  the  aluminium  family.  Yet  it  is  a  pronounced 
non-metal,  and  its  oxide  and  hydroxide  are  almost  wholly  acidic ; 
aluminium  is  a  metal,  and  with  its  oxide  and  hydroxide  basic  proper- 
ties predominate.  Boron  and  its  compounds  really  resemble  carbon 
and  silicon  and  their  compounds  in  all  chemical  properties  except  the 
property  of  valence. 

Occurrence.  —  Like  silicon,  boron  is  found  in  oxygen  compounds, 
namely,  in  boric  acid  (q.v.)  and  its  salts.  Of  the  latter,  sodium  tetra- 
borate  Na2B407,  or  borax,  came  first  from  India  under  the  name  of  tin- 
cal.  It  constitutes  a  large  deposit  in  Borax  Lake  in  California.  Cole- 
manite,  Ca^BgO^  5H20,  from  California,  and  other  complex  borates, 
furnish  a  large  part  of  the  commercial  supply  of  compounds  of  boron. 

Preparation.  —  When  boric  oxide  is  heated  with  powdered  mag- 
nesium (B203  -f-  3Mg  — >  3MgO  -f  2B),  amorphous  boron  can  be  sep- 
arated with  some  difficulty  from  the  borides  of  magnesium  in  the 
resulting  mixture.  When  excess  of  powdered  aluminium  is  used,  hard 
crystals  of  boron,  containing  aluminium,  are  found  in  the  solidified 
metal.  They  may  be  separated  by  interaction  of  the  metal  with  dilute 
hydrochloric  acid. 

Properties.  —  Amorphous  boron  is  a  black  powder.  It  unites 
with  the  same  elements  as  does  silicon  (p.  519),  but  with  somewhat 
greater  activity.  Like  carbon  (p.  379),  it  is  also  oxidized  by  hot,  con- 


SILICON   AND   BORON  527 

centrated  sulphuric  or  nitric  acid,  the  product  being  boric  acid.  The 
crystalline  variety  is  less  rapidly  attacked  in  each  case.  Both  kinds 
interact  with  fused  potassium  hydroxide,  giving  a  borate : 

2B  +  6KOH  ->  2K8B08  +  3H2. 

Hydrides  and  Halides  of  Boron.  —  When  magnesium  boride 
Mg3B2  is  treated  with  hydrochloric  acid,  a-  gas  containing  much  hydro- 
gen, and  possibly  several  hydrides  of  boron,  is  given  off.  By  cooling 
the  mixture  with  liquid  air,  Eamsay  obtained  a  white  solid  which, 
when  it  resumed  the  gaseous  condition,  appeared  to  be  B3H8. 

By  combined  action  of  carbon  and  chlorine  on  boric  oxide,  using 
the  principle  employed  in  preparing  silicon  tetrachloride  (p.  520),  the 
trichloride  of  boron  may  be  made.  It  is  likewise  formed  by  direct 
union  of  the  free  elements.  It  is  also  made  easily  by  heating  boric 
acid  and  phosphorus  pentachloride,  the  action  being  an  example  of  the 
behavior  of  the  latter  towards  hydroxyl  compounds  (of.  p.  463) : 

B(OH)3  +  3PC16->BC13  +  3POC18  +  3HC1. 

The  products  are  separated  by  fractional  distillation.  Boron  trichloride 
is  a  liquid  which  boils  at  18Q,  fumes  strongly  in  moist  air,  and  is  com- 
pletely hydrolyzed  by  water. 

Boron  trifluoride  BF3  is  made  by  the  interaction  of  calcium  fluoride 
and  sulphuric  acid  with  boron  trioxide.  The  mode  of  preparation  and 
the  properties  of  the  substance  recall  silicon  tetrafluoride  (p.  521).  It 
interacts  with  water,  like  the  latter,  giving  boric  acid  and  hydrofluo- 
boric  acid : 

4BF8  +  3H20  ->  B(OH)8  +  3HBF4. 

The  boric  acid,  not  being  very  soluble,  is  precipitated.  Hydrofluoboric 
acid  is  known  only  in  solution,  although  many  of  its  salts  are  stable. 

Boric  Acid.  —  Boric  acid  (boracic  acid)  is  somewhat  volatile  with 
steam  (cf.  p.  471),  and  is  found  in  Tuscany  in  jets  of  water  vapor 
(soffioni)  which  issue  from  the  ground.  Water,  retained  in  small 
basins  by  brickwork,  is  placed  over  the  openings,  and  from  this  water, 
after  evaporation  by  the  help  of  the  steam  of  the  soffioni  themselves, 
boric  acid  is  obtained  in  crystalline  form.  As  boric  acid  is  very 
feeble,  and  withal  little  soluble,  it  may  also  be  made  by  interaction  of 
sulphuric  acid  and  concentrated  borax  solution : 

Na2B407  +  H2S04  +  5H2O  <=»  Na^SO,  +  4H8BOJ. 


528  INORGANIC   CHEMISTRY 

Boric  acid  crystallizes  from  water  in  thin  white  plates,  which  are 
soapy  (like  graphite  and  talc)  to  the  touch.  Its  solubility  in  water  is  4 
parts  in  100  at  19°  and  34  in  100  at  100°.  The  solution  scarcely  affects 
litmus.  It  confers  a  green  tint  on  the  Bunsen  flame.  The  effect  is 
best  seen  by  setting  fire  to  the  vapor  of  a  boiling  alcoholic  solution  of 
the  acid.  This  behavior  is  used  as  a  test  for  the  acid.  At  100°  the  acid 
slowly  loses  water,  leaving  metaboric  acid  HB02,  and  at  140°  tetra- 
boric  acid  is  formed  :  4HB02  —  H20  —  >  H2B407.  Strong  heating  gives 
the  oxide  B2O3.  When  dissolved  in  water,  these  dehydrated  compounds 
revert  to  boric  acid.  The  solution  of  boric  acid  in  water  is  used  as  an 
antiseptic  in  medicine,  and  as  a  preservative  for  milk  and  other  foods. 

Borates.  —  Borates  derived  from  orthoboric  acid  are  practically 
unknown.  The  most  familiar  salt  is  borax  or  sodium  tetraborate. 
The  decahydrate  Na2B407,  10H20,  which  crystallizes  from  water  at  27° 
in  large,  transparent  prisms,  and  the  pentahydrate  which  crystallizes 
at  56°,*  are  both  marketed.  They  are  made  by  crystallization  of  native 
borax.  In  Germany,  borax  is  prepared  from  boracite,  found  at  Stass- 
furt,  by  decomposing  a  solution  of  the  mineral  with  hydrochloric  acid  : 

MgCl2,  2Mg3B8015  +  12HC1  +  18H2O  ->  7MgCl2  +  16B(OH)8. 

The  boric  acid  is  redissolved  in  boiling  water,  and  sodium  carbonate 
is  added  : 


4B(OH)8  +  NasCOa-^NaAO,  +  6H2O  +  C02. 

In  California  it  is  made  from  colemanite  by  interaction  with  sodium 
carbonate. 

Since  boric  acid  is  a  feeble  acid,  borax  is  extensively  hydrolyzed  by 
water,  and  the  solution  has  a  marked  alkaline  reaction. 

When  heated  with  oxides  of  metals,  sodium  tetraborate  behaves 
like  sodium  metaphosphate  (cf.  p.  468),  and  is  used  in  the  form  of  beads 
in  analysis.  If  its  formula  be  written  2NaB02,B208,  it  will  be  seen 
that  a  considerable  excess  of  the  acid  anhydride  is  contained  in  it,  and 
that,  therefore,  a  mixed  metaborate  may  be  formed  by  union  with 
some  basic  oxide.  Thus,  with  a  trace  of  cupric  oxide,  the  bead  is 
tinged  with  green,  from  the  presence  of  a  compound  like  2NaB02, 
Cu(BO2)2.  Cobalt  compounds  give  a  deep-blue  color  to  the  bead.  For 
the  same  reason,  borax  is  used  in  hard  soldering.  The  hard  solder 

*  For  explanation  of  the  relation  of  temperature  of  crystallization  to  degree  of 
hydration,  see  under  Manganous  sulphate. 


SILICON   AND   BORON  529 

(brass)  is  placed,  with  a  little  borax,  upon  the  joint  between  the  objects 
of  copper  or  brass  which  are  to  be  soldered.  At  the  temperature  pro- 
duced by  the  blast-lamp  the  borax  dissolves  the  superficial  coating  of 
oxides,  and  the  molten  solder  is  able  to  "  wet "  the  clean  surfaces.  A 
substance  used  thus,  to  bring  infusible  bodies  into  a  fusible  form  of 
combination,  is  called  a  flux. 

Boron  Trioxide.  —  The  oxide,  as  made  by  heating  boric  acid,  is  a 
glassy  white  solid.  It  is  obtained  also  by  burning  boron  in  oxygen. 
Being  almost  perfectly  involatile,  it  is  able,  when  heated  with  salts,  to 
displace  other  acid  anhydrides  which  can  be  vaporized  : 

K2S04  +  B2O8  <=±  2KB02  +  S03f . 

It  has  a  slight  tendency  to  act  as  a  basic  oxide.  With  fuming  sul- 
phuric acid  it  gives  a  boryl  pyrosulphate  BO,HS207  which  is  decom- 
posed by  water,  and  with  phosphoric  acid  a  phosphate  BP04  which  is 
stable. 

Nitride  and  Carbide.  —  One  of  the  difficulties  in  making  free 
boron  is  due  to  its  very  great  affinity  for  nitrogen,  with  which,  when 
heated  in  the  air,  it  unites  to  form  a  nitride  BN.  This  compound  is 
more  easily  made  by  heating  borax  with  ammonium  chloride  : 

NaaB407  -f  4NH4C1  ->  4BN  +  2NaCl  +  7H20  -f  2HC1. 

The  nitride  is  a  white  solid  which  is  easily  decomposed  when  heated 
in  a  current  of  steam  (cf.  p.  417) : 

BN  +  3H20  ->  B(OH)3  +  NH3. 

A  carbide  of  boron  B6C  is  made  by  heating  the  free  substances  in 
the  electric  furnace.  It  is  harder  than  carborundum,,  and  stands 
next  to  the  diamond  in  respect  to  hardness. 

Exercises.  —  1.  Why  is.  the  fact  that  carborundum  is  not  affected 
by  water  or  acids  worthy  of  mention  ? 

2.  Compare  and  contrast  the   elements    carbon   and   silicon,  and 
their  corresponding  compounds. 

3.  Why  are  there  no  colloidal  solutions  of  iron  and  zinc  (p.  524)  ? 

4.  What  would  be  the  interaction  between  aqueous  solutions  of  an 
ammonium  salt  and  of  sodium  orthosilicate  ? 


CHAPTER    XXXII 
THE   BASE-FORMING   ELEMENTS 

THERE  are  two  ways  in  which  the  chemistry  of  a  given  set  of  ele- 
ments may  be  described.  We  may  take  up  the  elements  in  succession, 
and  discuss  under  each  its  physical  and  chemical  properties,  and  the 
manufacture  and  behavior  of  a  certain  number  of  its  compounds,  such 
as  the  oxide,  hydroxide,  nitrate,  and  sulphate.  Or  we  may  arrange  our 
major  classification  according  to  the  properties  and  the  forms  of  com- 
bination, and  detail  the  facts  about  the  same  set  of  elements  under 
each.  Both  methods  have  such  advantages  that  neither  can  be  sacri- 
ficed entirely.  We  shall,  therefore,  adopt  the  former  plan  for  our  divis- 
ion into  chapters,  following  the  usage  already  employed  for  the  non- 
metals.  In  the  present  chapter  a  preliminary  view  of  the  chemistry 
of  the  metals  will  be  given  according  to  the  second  method. 

Physical  Properties  of  the  Metals.  —  A  knowledge  of  the  phys- 
ical properties  of  the  metals  is,  to  the  chemist,  of  the  greatest  impor- 
tance in  connection  with  their  manufacture  and  treatment.  The 
following  brief  statement  in  regard  to  some  of  these  properties  is  illus- 
trative rather  than  exhaustive.  It  should  be  noticed  that  the  properties 
of  a  metal  vary  according  as  the  specimen  has  been  prepared  by  roll- 
ing, casting,  or  some  other  process.  Numerical  values,  therefore,  when 
given,  are  only  approximate. 

Metals  show  what  is  commonly  called  a  metallic  luster,  but,  as  a 
rule,  they  do  so  only  when  in  compact  form.  Magnesium  and  alumin- 
ium exhibit  it  when  powdered,  but  most  of  the  metals  when  in  this 
condition  are  black.  In  compact  masses  the  metals  are  usually  silvery 
white  in  color.  Gold  and  copper,  which  are  yellow  and  red  respect- 
ively, are  the  conspicuous  exceptions. 

The  metals  can  all  be  obtained  in  crystallized  form,  when  a  fused 
mass  is  allowed  to  cool  slowly  and  the  unsolidified  portion  is  poured 
off.  In  almost  all  cases  the  crystals  belong  to  the  regular  system. 
With  the  metals  most  nearly  allied  to  the  non-metals,  however,  they  do 
not.  Thus,  the  crystals  of  antimony  and  bismuth  belong  to  the  hex- 
agonal system,  and  those  of  tin  to  the  square  prismatic. 

530 


THE   BASE-FORMING   ELEMENTS 


531 


The  metals  vary  in  specific  gravity  from  lithium,  which  is  little 
more  than  half  as  heavy  as  water  (sp.  gr.  0.53),  to  osmium,  whose 
specific  gravity  is  22.5.  Those  which  have  a  specific  gravity  less  than 
5,  namely,  potassium,  sodium,  calcium,  magnesium,  aluminium,  and 
barium,  are  called  the  light  metals,  and  the  others  the  heavy  metals. 

Most  metals  are  malleable,  and  can  be  beaten  into  thin  sheets  with- 
out loss  of  continuity.  Those  which  are  allied  to  the  non-metals,  how- 
ever, such  as  arsenic,  antimony,  and  bismuth,  are  brittle,  and  can 
be  reduced  to  powder  in  a  mortar.  Zinc  becomes  malleable  only  when 
heated  to  150°.  The  order  of  the  elements  in  respect  to  this  property, 
beginning  with  the  most  malleable,  is :  Au,  Ag,  Cu,  Sn,  Pt,  Pb,  Zn, 
Fe,  Ni. 

The  tenacity  of  the  metals  places  them  in  an  order  different  from 
the  above.  It  is  measured  by  the  number  of  kilograms  which  a  piece 
of  the  metal  1  sq.  mm.  in  section  can  sustain  without  breaking.  The 
values  are  as  follows  :  Fe  62,  Cu  42,  Pt  34,  Ag  29,  Au  27,  Al  20,  Zn  5, 
Pb2. 

The  hardness  is  measured  by  the  ease  with  which  the  material  may 
be  disintegrated  by  a  sharp,  hard  instrument.  Potassium  is  as  soft  as 
wax,  while  chromium  is  hard  enough  to  cut  glass. 

The  temperature  at  which  the  metal  fuses  has  an  important  bearing 
on  its  manufacture.  Most  of  the  following  melting-points  are  only 
approximate  : 


Mercury   . 

-40° 

Zinc  . 

420° 

Cast  iron  . 

1160° 

Potassium 

62° 

Antimony 

437° 

Nickel.     . 

1500° 

Sodium     . 

96° 

Aluminium 

700°  (?) 

Platinum  . 

1780° 

Tin 

230° 

Magnesium 

750° 

Iron  (pure) 

1800° 

Bismuth   . 

264° 

Silver      . 

954° 

Manganese 

1900° 

Cadmium 

320° 

Copper    . 

1057° 

Chromium 

2000° 

Lead   .      . 

326° 

Gold  .      . 

1075° 

Iridium    . 

2200° 

It  will  be  seen  that  mercury  is  a  liquid,  that  potassium  and  sodium 
melt  below  the  boiling-point  of  water,  and  that  the  metals  down  to  the 
foot  of  the  second  column  can  be  melted  easily  with  the  Bunsen  flame. 
The  metals  osmium,  molybdenum,  uranium,  tungsten,  and  vanadium 
have  melting-points  above  the  melting-point  of  platinum. 

The  methods  of  manufacture  and  the  treatment  of  metals  are  much 
influenced  also  by  their  volatility.  The  following  are  easily  distilled : 
Mercury,  b.-p.  357° ;  potassium  and  sodium,  b.-p.  about  700°  ;  cadmium, 
b.-p.  770° ;  zinc,  b.-p.  950°.  Even  the  most  in  volatile  metals  can  be 
converted  into  vapor  in  the  electric  arc. 


532  INORGANIC   CHEMISTRY 

In  many  cases  molten  metals  dissolve  in  one  another  freely.  Tli6 
mixtures  are  called  alloys,  and  in  some  cases  have  the  properties  of 
solid  solutions.  Sometimes,  as  in  the  case  of  lead  and  tin,  mixtures  can 
be  formed  in  all  proportions.  On  the  other  hand,  the  solubility  may  be 
limited,  as  in  the  case  of  zinc  and  lead,  where  only  1.6  parts  of  the 
former  dissolve  in  100  parts  of  the  latter.  The  colors  of  alloys  are 
not  the  average  of  those  of  the  constituents.  Thus,  a  mixture  contain- 
ing copper  with  30  per  cent  of  tin  is  perfectly  white.  A  similar  mix- 
ture with  30  per  cent  of  zinc  is  pale  yellow.  The  nickel  alloy  used  in 
coining  contains  75  per  cent  of  copper  and  25  per  cent  of  nickel,  yet  it 
shows  none  of  the  color  of  the  former.  Thirty  per  cent  of  gold  may 
be  added  to  silver  without  conferring  any  yellow  tint  upon  it. 

Some  of  the  properties  of  alloys  are  classifiable  by  the  ordinary 
laws  of  solution.  Thus,  a  foreign  body  lowers  the  vapor  tension  of  a 
solvent  (p.  161),  and  so  the  presence  of  a  foreign  metal  diminishes  the 
ease  with  which  particles  can  be  torn  from  the  surface  by  any  means 
whatever.  That  is  to  say,  it  increases  the  hardness.  A  foreign  metal 
also  lowers  the  melting-point  (p.  163).  In  many  cases  the  metal  be- 
comes less  active  when  alloyed.  Thus,  a  mixture  of  gold  and  silver 
containing  twenty-five  per  cent  of  the  latter  does  not  interact  visibly 
with  nitric  acid.  It  is  necessary  to  bring  the  amount  of  silver  up  to 
75  per  cent  at  least  ("quartation")  in  order  that  the  silver  may  be 
freely  attacked  by  the  warm  acid.  The  gold  remains  in  any  case  un- 
touched. The  malleability  and  the  conductivity  for  heat  and  elec- 
tricity are  diminished  by  addition  of  a  foreign  metal.  Copper,  whose 
commercial  applications  depend  largely  on  the  first  and  third  of  these 
properties,  is  much  affected  in  respect  to  them  by  the  presence  of  even 
small  traces  of  impurities. 

Alloys  in  which  mercury  forms  one  of  the  components  are  known 
as  amalgams  (Gk.  /taXay/wi,  a  soft  mass),  and  are  formed  with  especial 
ease  by  the  lighter  metals.  Of  the  common  metals,  iron  is  the  least 
miscible  with  mercury. 

The  good  conductivity  of  metals  for  electricity  distinguishes  them 
with  some  degree  of  sharpness  from  the  non-metals.  They  show  con- 
siderable variation  amongst  themselves,  silver  conducting  sixty  times 
as  well  as  mercury.  The  conductivity  increases  as  the  temperature  is 
lowered,  and  this  fact  is  taken  advantage  of  in  the  measurement  of  the 
temperature  of  liquefied  gases.  The  platinum  resistance  thermometer 
consists  chiefly  of  a  wire  of  platinum.  The  resistance  of  this  metal 
diminishes  so  rapidly,  with  decreasing  temperature,  that  measurement 


THE   BASE-FORMING  ELEMENTS  533 

of  its  resistance  can  be  used  for  the  accurate  determination  of  the  tem- 
perature. In  the  following  table  the  conductivities  of  the  metals  are 
expressed  in  terms  of  the  number  of  meters  of  wire  1  sq.  mm.  in  sec- 
tion which,  at  15°,  offer  a  resistance  of  one  ohm : 

Silver,  cast   .      .      .      .62.89  Nickel,  cast  .      .      .  7.59 

Copper,  commercial      .   57.40  Iron,  drawn        .      .7.55 

Gold,  cast     .      .      .  •    .  46.30  Platinum       .      .  5.7-8.4 

Aluminium,  commercial  31.52  Steel 5.43 

Zinc,  rolled  ....  16.95  Lead        ....  4.56 

Brass 14.17  Mercury  ....  1.049 

The  resistance  at  0°  of  a  column  of  mercury  1  sq.  mm.  in  section 
and  1.063  meters  long  is  called  the  ohm,  and  is  employed  in  express- 
ing the  conductivities  of  solutions  (p.  328).  To  compare  the  above  fig- 
ures with  those  given  for  solutions,  however,  it  must  be  recalled  that, 
in  the  measurement  of  the  conductivities  of  the  latter,  a  column  only  1 
cm.  in  length  and  of  1  sq.  cm.  area  was  employed,  so  that  the  figures 
representing  the  conductivities  of  solutions  are  on  a  scale  approxi- 
mately ten  thousand  times  as  great  as  those  presented  in  the  above 
table.  Thus,  normal  hydrochloric  acid  (p.  328)  has  a  conductivity  on 
the  above  scale  of  0.0301,  or  less  than  a  thirtieth  of  that  of  mercury. 


General    Chemical  Relations  of  the   Metallic   Elements.— 

Since  most  of  the  compounds  of  the  metals  are  ionogens,  their  solutions, 
except  when  the  metal  is  a  part  of  a  compound  ion  or  of  a  complex  ion 
(see  below),  all  contain  the  metal  in  the  ionic  state,  and  the  resulting 
substances,  such  as  kalion  and  cuprion,  have  constant  properties, 
irrespective  of  the  nature  of  the  negative  ion  with  which  they  may  be 
mixed.  The  properties  of  the  ions,  simple  and  compound,  are  much 
used  in  making  tests  in  analytical  chemistry.  On  the  other  hand,  the 
chemical  properties  of  the  oxides  and  of  the  salts  in  the  dry  state  are 
of  importance  in  connection  with  metallurgy. 

There  are  six  chemical  properties  which  are  more  or  less  generally 
characteristic  of  the  metallic  elements.  The  first  two  of  them  have 
already  been  discussed  somewhat  fully. 

1.  The  metals  are  able  by  themselves  to  form  positive  radicals  of 
salts,  and,  therefore,  to  exist  alone  as  positive  ions  (pp.  337,  404). 

2.  The  oxides  and  hydroxides  of  the  metals  are  basic  (pp.  119,  404). 

3.  The  halogen  compounds  of  the  typical  metals  are  little,  if  at  all, 
hydrolyzed  by  water  (see  next  section). 


534  INORGANIC  CHEMISTRY 

These  three  properties  are  characteristic  of  all  metals,  and  form  the 
basis  of  the  distinction  between  metals  and  non-metals.  The  remaining 
three  apply  to  many  metals,  but  no  one  of  them  applies  to  all. 

4.  An  oxide  or  hydroxide  which  is  basic  may  also  be  acidic,  as,  for 
example,  zinc  hydroxide  (p.  405).     Even  when  this  is  not  the  case, 
some  other  oxide  of  the  metal  may  be  acidic  exclusively,  as  is  manga- 
nese heptoxide  (p.  405).     In  consequence  of  either  of  these  facts,  a 
metal  may  form  part  of  the  negative  radical  of  a  simple  salt,  and  there- 
fore be  found  in  a  negative  ion,  as,  Zn02"  or  MnO/. 

5.  Some  salts  of  certain  metals  combine  with  those  of  others  to 
give  complex  salts  (p.  363  and  next  section  but  one).    Of  this  sort  are 
the  complex  cyanides,  such  as  K.Ag(CISr)2  and  K4.Fe(CN)6.     A  metal 
thus  forms  part  of  the  negative  radical  of  a  salt  of  a  complex  acid, 
and  therefore  is  found  in  an  anion  like  Ag(CN)2'  or  PtCl6". 

6.  Some  metals  also  form  parts  of  complex  cations  which  are  con- 
tained in  solutions  of  molecular  compounds  (p.  443),  chiefly  those  of 
salts  with  ammonia.     Thus,  when  AgCl,  3NH8  is  dissolved  in  water,  or 
when  ammonium  hydroxide  is  added  to  a  solution  of  a  salt  of  silver,  the 
positive  ion  is  found  to  be  Ag(NH8)2*  (see  under  Copper).      (For  a  de- 
tailed illustration  of  the  application  of  these  six  criteria,  see  discussion 
of  the  chemical  relations  in  the  nitrogen  family,  Chap,  xli.) 

Aside  from  these  points,  many  features  in  the  behavior  of  metals 
and  their  compounds  are  summed  up  in  the  electromotive  series  (p.  362). 
The  reader  should  re-read  all  the  parts  referred  to  above  before  proceed- 
ing farther.  He  should  also  reexamine  the  various  kinds  of  chemical 
changes  discussed  on  p.  187  and  particularly  the  varieties  of  ionic 
chemical  change  on  p.  360. 

• 

Hydrolysis  of  Halogen  Compounds,  Used  to  Distinguish 
Metals  from  Non-Metals.  —  We  have  seen  that  the  halogen  com- 
pounds of  phosphorus  (p.  463),  of  sulphur  (p.  398),  of  silicon  (p.  521), 
and  of  other  non-metals,  are  completely  hydrolyzed  by  water,  giving 
the  hydrogen  halide,  and  an  acid  which  contains  the  hydroxyl  of  the 
water : 

PC18  +  3HOH  -» 3HC1  +  P(OH)8. 

Now,  those  elements  whose  halogen  compounds  are  not  hydrolyzed  by 
water,  or,  at  all  events,  are  only  partly  hydrolyzed,  are  the  ones  classed 
as  metals.  Thus,  sodium  chloride  is  not  decomposed  appreciably  by 


THE  BASE-FORMING  ELEMENTS  535 

water,  and  cupric  chloride,  like  cupric  sulphate  (p.  344),  is  but  slightly 
hydrolyzed,  and  its  solution  has  a  faint  acid  reaction  : 

CuCl2  +  2H2O  ±3  2HC1  +  Cu(OH)2. 

In  a  few  cases,  as  with  the  chlorides  of  antimony  (q.v.)  and  of  bis- 
muth (q.v.)}  a  considerable  proportion,  but  not  all,  of  the  halogen  is  re- 
moved from  each  molecule  : 

SbCl8  +  H20^±2HC1  +  SbOClJ. 


The  resulting  mixture  is  strongly  acid,  and  the  product  (antimony 
oxy  chloride)  is  a  definite  compound,  of  the  nature  of  a  mixed  salt 
(p.  359),  known  as  a  basic  salt.  The  difference  is  that,  with  the  halides 
of  the  metals,  the  action  of  water  is  notably  reversible,  while  with  halides 
of  the  non-metals  it  is  not  so. 

Hydrolysis  of  the  halides  of  the  metals  is  increased  by  rise  in  tem- 
perature and  by  dilution  of  the  solution  (addition  of  more  water),  and 
also  gains  headway  when  one  of  the  products  of  hydrolysis  is  thrown 
down  as  a  precipitate.  The  last  two  influences  are  the  ones  which 
normally  permit  a  reversible  action  to  approach  completion  (p.  259). 

The  halogen  compounds  are  chosen  as  the  basis  of  this  criterion 
because  the  halogen  acids  are  active  and  would  reverse  the  hydrolysis 
completely,  and  leave  no  acid  reaction,  if  the  result  depended  upon 
them  alone.  It  is  the  lack  of  activity  in  the  base,  and  the  ten- 
dency of  its  molecules  to  be  formed  from  the  metal  ions  of  the  salt 
and  the  OH'  of  the  water  (p.  314),  that  determine  the  slight  hydrolysis, 
when  it  occurs.  Thus,  this  criterion  is  simply  another  means  of  recog- 
nizing whether  or  not  the  hydroxide  of  the  element  is  a  strong  (much 
ionized)  base,  and  its  application  gives,  therefore,  the  same  result  in 
each  case  as  does  the  employment  of  the  second  of  the  chemical  charac- 
teristics of  metallic  elements  (see  above). 

Other,  non-halide  salts  of  the  metals,  even  of  the  most  active,  may  be 
extensively  hydrolyzed  by  water.  Thus,  sodium  sulphide  is  decomposed 
by  it  (p.  375)  to  the  extent  of  one-half.  But  here  the  solution  is 
alkaline  in  reaction  : 

Na2S  -f  H.O  -»  NaSH  +  NaOH, 

owing  to  the  small  ionization  of  the  SH'  ion,  and  the  result  is  due  to 
the  feebleness  of  the  acid  (H2S).  Indeed,  the  great  activity  of  the 
base  is  demonstrated  by  the  final  reaction  of  the  solution,  and  the 


536  INORGANIC   CHEMISTRY 

metallic  nature  of  sodium  is  therefore  not  impugned  by  the  existence 
of  hydrolysis  per  se  in  such  a  salt  as  this,  but  is  rather  confirmed. 

To  sum  up  :  An  alkaline  reaction  shows  that  we  have  a  solution  of 
a  salt  of  an  active  metal  with  a  weak  acid  ;  an  acid  reaction  that  we 
have  a  salt  of  an  inactive  metal,  which  may  even  verge  on  the  non- 
metallic,  with  an  active  acid.  Salts  of  two  active  components  give 
neutral  solutions.  Thus,  as  a  particular  case,  the  halogen  compounds 
of  the  typical  metals  are  not  perceptibly  hydrolyzed  by  water,  and  the 
hydrolysis  of  the  halide  of  a  less  pronounced  metal  takes  place  with 
acid  reaction,  and  is  easily  reversible  by  excess  of  either  product.  A 
salt  of  an  acid  and  base  both  of  which  are  weak  is  hydrolyzed,  often 
completely.  Aluminium  carbonate  and  ammonium  silicate  are  examples 
of  salts  which,  for  this  reason,  are  completely  hydrolyzed.  The  result- 
ing mixture  may  have  an  acid  or  a  basic  reaction,  if  the  acid  or  the 
base  is  sufficiently  soluble  and  sufficiently  active.  Thus,  ammonium 
sulphide  solution  is  alkaline. 

Salts  of  Complex  Acids.  —  These  salts  are  of  many  classes,  and 
arise  by  direct  union  of  two  salts.  Thus  we  have  cyanides  like 
K.Ag(CN)2  (potassium  argenticyanide),  K.Cu(CN)2  (potassium  cupro- 
cyanide),  and  K4.Fe(CN)6  (potassium  ferrocyanide)  : 

K.CN  +  AgCN  f  ->K.Ag(CN)2. 

The  complex  sulphur  compounds  of  arsenic,  antimony,  and  tin  (y.v.), 
such  as  sodium  sulphostannate  Na2.SnS8,  are  made  in  the  same  way  : 

4-  SnS2  f 


Many  double  halides  are  of  a  like  nature,  as  K2.PtClg  (potassium  chloro- 
platinate),  H2.PtCl<j  (chloroplatinic  acid),  and  Na.AuCl4  (sodium 
chloraurate).  In  every  case  the  metal  of  one  of  the  original  salts  is 
contained  in  the  negative  radical  (the  anion).  Hydrofluosilicic  acid 
(p.  521)  is  a  compound  analogous  to  those  of  the  last  group,  excepting 
that  SiF^  from  which  it  is  formed,  is  not  a  salt,  and  silicon  is  not  a  metal. 
The  characteristic  of  these  compounds  is  that,  when  they  are 
ionized,  the  less  positive  metal  is  contained  in  a  complex  anion  like 
Ag(CN)2',  SnS8",  PtClg".  In  fact,  they  behave,  in  most  respects,  like 
ordinary  single  salts.  Thus,  they  undergo  double  decomposition  in  the 
normal  manner,  the  complex  ion  acting  as  a  whole.  For  example,  a 
soluble  ferrocyanide  with  a  zinc  salt  gives  zinc  ferrocyanide  : 

K4.Fe(CN)6  4-  2Zn.SO4^±  Zn2Fe(CN)6|  -f  2K2.S04. 


THE   BASE-FORMING   ELEMENTS  537 

This  behavior  distinguishes  salts  of  complex  acids  from  double  salts 
(p.  360).  In  typical  cases  the  latter  resemble  the  former,  indeed, 
only  in  their  mode  of  preparation  (by  union  of  two  simple  salts),  and 
in  solution  are  resolved  once  more  into  the  component  salts  and  their 
separate  ions.  The  distinction  must  be  regarded,  however,  only  as  a 
means  of  rough  classification  for  practical  purposes.  In  reality,  all 
complex  ions  give  at  least  a  trace  of  the  simpler  ions,  and.  many  are  de- 
composed to  a  noticeable  extent.  At  the  other  extreme,  the  double  salts 
form  complex  ions  in  appreciable  amounts  in  concentrated  solutions. 

In  harmony  with  the  above  characteristic,  salts  of  complex  acids, 
like  potassium  ferrocyanide,  when  treated  with  acids,  like  sulphuric 
acid,  undergo  double  decomposition,  and  give  the  free  complex  acid 
(here  ferrocyanic  acid)  : 

K4.Fe(CN)6  +  2H2.S04?=»H4Fe(CN)6j  +  2K2.S04. 

But  in  many  cases  the  acids  are  unstable,  those  of  the  cyanides,  for 
example,  giving  up  hydrocyanic  acid  and  those  of  the  complex  sulphur 
compounds,  hydrogen  sulphide  : 


The  complex  halogen  acids,  however,  like  H^PtClg  and  H.AuCl4,  are 
stable,  and  are  fairly  active  as  acids. 

The  similarity  of  these  compounds  to  oxygen  acids,  and  their  salts,  may  be  seen 
if  we  imagine  them  to  be  cases  in  which  cyanogen,  sulphur,  chlorine,  and  other 
radicals  have  taken  the  place  of  oxygen  in  the  anion.  Thus,  [(CN)2]n,  [C1J11,  and 
S"  are  equivalent  to  On.  Many  of  the  corresponding  oxygen  compounds  are 
actually  known,  as  sodium  stannate  Na2.SnO3  and  sodium  aurate  Na.AuO2,  corre- 
sponding respectively  to  Na2.SnS3  and  Na.AuCl4. 

The  behavior  of  complex  ions  is  discussed  further  under  Copper. 

Classification  of  the  Metallic  Elements  by  their  Chemical 
Relations.  —  In  treating  of  the  metallic  elements  and  their  com- 
pounds we  shall  use  the  groupings  provided  by  the  periodic  system 
(p.  411).  A  division  into  eleven  sets,  which  are  described  briefly,  and 
in  general  terms  below,  will  be  sufficient  for  the  purpose  of  this  book  : 

1.  Metals  of    the  Alkalies.  —  Lithium,  sodium,    potassium,    rubi- 
dium, caesium,  and  the  radical   ammonium  NH4.     These  metals  are 
univalent,  and  their  oxides  and  hydroxides  have  strongly  basic  proper- 
ties.    Their  salts  with  active  acids  are  not  hydrolyzed  in  solution. 

2.  Metals  of  the  Alkaline  Earths.  —  Calcium,  strontium,  barium, 


538  INORGANIC   CHEMISTRY 

and,  possibly,  radium.  These  metals  are  bivalent  in  all  their  com- 
pounds. Their  oxides  and  hydroxides  are  strongly  basic,  but  the 
latter  are  not  so  soluble  in  water  as  are  the  hydroxides  of  the  former 
family.  The  salts  with  active  acids  are  not  hydrolyzed. 

3.  Copper,  Silver,  and  Gold.  —  These    metals    occupy   the  right- 
hand  side  of  the  second  column  of  the  table  (p.  411).     Their  alliance 
with  the  alkali  metals,  their   neighbors,  is  rather  remote.     Each  of 
them,  however,  gives  compounds  in  which  it  is  univalent,  although,  in 
their  commoner  compounds,  copper  and  gold  are  bivalent  and  triva- 
lent,  respectively.     The  oxides  and  hydroxides  of  copper  and  gold  have 
rather  weak  basic  properties ;  those  of  silver  are  much  more  active. 

4.  Beryllium,  Magnesium,  Zinc,  Cadmium,  and  Mercury.  —  These 
metals,  occupying  the  right-hand  side  of  the  third  column,  are  biva- 
lent, although  mercury  forms  a  series  of  compounds  in  which  it  is 
univalent  as  well.     The  oxides  and  hydroxides  are  feebly  basic,  those 
of  magnesium  being  the  most  basic. 

5.  Aluminium  and  the  other  metals  on  both  sides  of  the  fourth 
column.  —  The  metals  of  these  groups  are  trivalent,  and  the  oxides  and 
hydroxides  are  feebly  basic  in  character.     The  hydroxide  of  aluminium 
has  also  a  feebly  acidic  tendency. 

6.  Germanium,  Tin,  and  Lead,  and  the  Titanium  Family.  —  In  ac- 
cordance with  their  position  in  the  periodic  table  these  metals  are  all 
quadrivalent.    At  the  same  time  they  act  also  as  bivalent  elements, 
the  compounds  of  this  class  in  the  case  of  lead  being  the  more  familiar. 
The  oxides  and  hydroxides  are  feebly  basic,  and  are  able  also  to  play 
the  role  of  acidic  oxides  towards  strong  bases. 

7.  Arsenic,  Antimony,  and  Bismuth,  and  the  metals  of  the  Vanadium 
Group.  —  These  elements,  like  nitrogen  and  phosphorus,  form  two  sets 
of  compounds  in  which  they  are  trivalent  and  quinquivalent,  respec- 
tively.    The  acidic  tendency  of  the  oxides  and  hydroxides,  which  in 
the  last  column  was  noticeable,  is  here  much  more  pronounced.    Both 
oxides  of  arsenic  are  almost  wholly  acidic  in  behavior,  and  the  pent- 
oxides  of  the  other  elements  are  likewise  acid  anhydrides  exclusively. 
The  trioxides  are  basic  in  a  feeble  way,  and  their  salts  are  much  hy- 
drolyzed by  water. 

8.  Chromium,  Molybdenum,  Tungsten,  and  Uranium.  —  These  ele- 
ments, occupying  the  left-hand  side  of  the  seventh  column,  exhibit  a 
considerable  variety  of  valence.     The  maximum,  however,  is  six  in  each 
case.    The  oxides  of  the  form  CrO  and  Cr203,  in  which  the  elements  are 
bivalent  and  trivalent,  are  base-forming.     Those  of  the  form  Cr08,  in 


THE   BASE-FORMING  ELEMENTS  539 

which  the  elements  are  sexivalent,  resemble  sulphur  trioxide  and  are 
acid  anhydrides. 

9.  Manganese.  —  This  element,  the  only  one  in  the  eighth  column 
which  has  not  yet  been  treated,  gives  several  series    of   compounds 
in   which   its   valence    varies    from    2    to    7.     Compounds    derived 
from  the   basic   oxide   MnO   are   the   salts   in  which   manganese   is 
most  distinctly  a  metallic  element.     The  highest  oxide,  Mn2O7,  is  an 
acid  anhydride. 

10.  Iron,  Nickel,  and  Cobalt.  —  These  elements  give  oxides  which 
are  feebly  basic.     Iron  gives  two  extensive  series  of   compounds  in 
which  it  is  bivalent  and  trivaleiit,  respectively.     Those  of  the  former  set 
resemble  the  bivalent  salts  of  manganese  and  zinc.     Those  of  the  latter 
resemble  the  salts  of  aluminium.     Cobalt  and  nickel  in  most  of  their 
compounds  are  bivalent  elements,  and  the  behavior  recalls  that  of  the 
compounds  of  bivalent  manganese  and  zinc. 

11.  Palladium  and  Platinum  Families. — The  metals  of  these  fam- 
ilies have  little  chemical  activity,  and  their  compounds  are  easily  decom- 
posed by  heating.     Along  with  gold,  silver,  and  mercury,  which  have 
similar  characteristics,  they  are  sometimes  grouped  together  under  the 
name  of  the  noble  metals. 

Occurrence  of  the  Metals  in  Nature. — The  minerals  from  which 
metals  are  extracted  are  known  as  ores.  They  present  a  compara- 
tively small  number  of  different  kinds  of  compounds.  Most  of  the 
metals  are  found  in  more  than  one  of  these  forms,  so  that  in  the  fol- 
lowing statement  the  same  metal  frequently  occurs  more  than  once. 

When  the  metal  occurs  free  in  nature  it  is  said  to  be  native.  Thus 
we  have  native  gold,  silver,  metals  of  the  platinum  group,  copper, 
mercury,  bismuth,  antimony,  and  arsenic  (cf.  p.  362). 

The  metals  whose  oxides  are  important  minerals  are  iron,  man- 
ganese, tin,  zinc,  copper,  and  aluminium.  The  metals  are  obtained 
commercially  from  the  oxides  in  each  of  these  cases. 

The  metals  whose  sulphides  are  used  as  ores  are  nickel,  cobalt, 
antimony,  lead,  cadmium,  zinc,  and  copper. 

From  the  carbonates  we  obtain  iron,  lead,  zinc,  and  copper.  Sev- 
eral other  metals,  such  as  manganese,  magnesium,  barium,  strontium, 
and  calcium  occur  in  larger  or  smaller  quantities  in  the  same  form  of 
combination. 

The  metals  which  occur  as  sulphates  are  those  whose  sulphates  are 
not  freely  soluble,  namely,  lead,  barium,  strontium,  and  calcium. 


540  INORGANIC   CHEMISTRY 

Compounds  of  metals  with,  the  halogens  are  not  so  numerous. 
Silver  chloride  furnishes  a  limited  amount  of  silver.  Sodium  and  pot- 
assium chlorides  are  found  in  the  salt-beds,  and  cryolite  3NaF,  A1F8  is 
used  in  the  manufacture  of  aluminium. 

The  natural  silicates  are  very  numerous,  but  are  seldom  used  for 
the  preparation  of  the  metals.  Many  of  them  are  employed  for  other 
commercial  purposes,  kaolin  (p.  525)  being  a  conspicuous  example  of 
this  class. 

Methods  of  Extraction  from  the  Ores.  —  The  art  of  extracting 
metals  from  their  ores  is  called  metallurgy.  Where  the  metal  is  native, 
the  process  is  simple,  since  melting  aw  ay  from  the  matrix  (p.  367)  is  all 
that  is  required.  Frequently  a  flux  (p.  529)  is  added,  which  combines 
with  the  matrix,  giving  a  fusible  slag.  Since  the  slag  is  a  melted  salt, 
usually  a  silicate,  and  does  not  mix  at  all  with  the  molten  metal,  sepa- 
ration of  the  products  is  easily  effected.  When  the  ore  is  a  compound, 
the  metal  has  to  be  liberated  by  furnishing  a  material  capable  of  com- 
bining with  the  other  constituent.  The  details  of  the  process  depend 
on  various  circumstances.  Thus  the  volatile  metals,  like  zinc  and 
mercury,  are  driven  off  in  the  form  of  vapor,  and  secured  by  conden- 
sation. The  involatile  metals,  like  copper  and  iron,  run  to  the  bottom  of 
the  furnace  and  are  tapped  off. 

Where  the  ore  is  an  oxide  it  is  usually  reduced  by  heating  with 
carbon  in  some  form.  This  holds  for  the  oxides  of  iron  and  copper, 
for  example.  Some  oxides  are  not  reducible  by  carbon  in  an  ordinary 
furnace.  Such  are  the  oxides  of  calcium,  strontium,  barium,  magnesium, 
aluminium,  and  the  members  of  the  chromium  group.  At  the  tempera- 
ture of  the  electric  furnace  even  these  may  be  reduced,  but  the  carbides 
are  formed  under  such  circumstances,  and  the  metals  are  more  easily 
obtained  otherwise.  Recently,  heating  the  pulverized  oxide  with  finely 
powdered  aluminium  has  come  into  use,  particularly  for  operations  on 
a  small  scale.  Iron  oxide  is  easily  reduced  by  this  means,  and  even 
the  metals  manganese  and  chromium  may  be  liberated  from  their 
oxides  quite  readily  by  this  action.  This  procedure  has  received  the 
name  aluminothermy  (g.v.)  on  account  of  the  great  amounts  of  heat 
liberated.  In  the  laboratory  the  oxides  of  the  less  active  metals  are 
frequently  reduced  in  a  stream  of  hydrogen  (cf.  p.  362). 

When  the  ore  is  a  carbonate,  it  is  first  heated  strongly  to  drive  out 
the  carbon  dioxide  (of.  p.  480)  :  FeC03<=±FeO  +  C02f ,  and  then  the 
oxide  is  treated  according  to  one  of  the  above  methods.  When  the 


THE   BASE-FORMING  ELEMENTS  541 

ore  is  a  sulphide,  it  has  to  be  roasted  (cf.  p.  378)  in  order  to  remove  the 
sulphur,  and  the  resulting  oxide  is  then  treated  as  described  above. 

Chlorides  and  fluorides  of  the  metals  can  be  decomposed  by  heat- 
ing with  metallic  sodium  (cf.  p.  518).  This  method  was  formerly  em- 
ployed in  the  making  of  magnesium  and  aluminium. 

The  metals  which  are  not  readily  secured  in  any  of  the  above  ways, 
can  be  obtained  easily  by  electrolysis  of  the. fused  chloride  or  of  some 
other  simple  compound.  Aluminium  is  now  manufactured  entirely  by 
the  electrolysis  of  a  solution  of  aluminium  oxide  in  molten  cryolite. 

Compounds  of  the  Metals  :  Oxides  and  Hydroxides.  —  The 
oxides  may  be  made  by  direct  burning  of  the  metal,  by  heating  the 
nitrates  (cf.  p.  444),  the  carbonates  (cf.  p.  480),  or  the  hydroxides  : 
Ca(OH)2<=>CaO  -+-  H2O  J.  They  are  practically  insoluble  in  water, 
although  those  of  the  metals  of  the  alkalies  and  of  the  metals  of  the  alka- 
line earths  interact  with  water  rapidly  to  give  the  hydroxides.  They  are 
usually  stable.  Those  of  gold,  platinum,  silver,  and  mercury  decompose 
when  heated,  yet  with  increasing  difficulty  in  this  order.  The  metals, 
like  the  non-metals,  frequently  give  several  different  oxides.  Those  of 
the  univalent  metals,  having  the  form  K20,  if  we  leave  cuprous  oxide 
and  aurous  oxide  out  of  account,  have  the  most  strongly  basic  qualities. 
Those  of  the  bivalent  metals  of  the  form  MgO,  when  this  is  the  only 
oxide  which  they  furnish,  are  base-forming.  Those  of  the  trivalent 
metals  of  the  form  ALjOg  are  the  least  basic  of  the  basic  oxides.  The 
oxides  of  the  forms  Sn02,  Sb205,  Cr03,  and  Mn2O7,  in  which  the  metals 
have  valences  from  4  to  7,  are  mainly  acid-forming  oxides,  although 
the  same  elements  usually  have  other  lower  oxides,  which  are  basic. 

The  hydroxides  are  formed,  in  the  cases  of  the  metals  of  the  alkalies 
and  alkaline  earths  by  direct  union  of  water  with  the  oxides.  They  are 
produced  also  by  double  decomposition  when  a  soluble  hydroxide  acts 
upon  a  salt  (cf.  p.  350).  All  hydroxides,  except  those  of  the  alkali  metals> 
lose  the  elements  of  water  when  heated,  and  the  oxide  remains.  In 
some  cases  the  loss  takes  place  by  stages,  just  as  was  the  case  witL 
orthophosphoric  acid  (p.  465).  Thus  lead  hydroxide  Pb(OH)2  (q.v.) 
first  gives  the  hydroxide  Pb20(OH)2,  then  Pb302(OH)2,  and  then  the 
oxide  PbO.  With  the  exception  of  those  of  the  metals  of  the  alkalies 
and  alkaline  earths,  all  the  hydroxides  are  little  soluble  in  water.  The 
hydroxides  of  mercury  and  silver,  if  they  are  formed  at  all,  are  evi- 
dently unstable,  for,  when  either  material  is  dried,  it  is  found  to  con- 
tain nothing  but  the  corresponding  oxide. 


542  INORGANIC  CHEMISTRY 

Compounds  of  the  Metals  :  Salts.  —  It  may  be  said,  in  general, 
that  each  metal  may  form  a  salt  by  combination  with  each  one  of 
the  acid  radicals.  In  the  succeeding  chapters  we  shall  describe  only 
those  salts  which  are  manufactured  commercially,  or  are  of  special  in- 
terest for  some  other  reason.  The  various  salts  will  be  described  under 
each  metal.  Here,  however,  a  few  remarks  may  be  made  about  the 
characteristics  of  the  more  common  groups  of  salts.  The  salts  are 
classified  according  to  the  acid  radicals  which  they  contain. 

The  chlorides  may  be  made  by  the  direct  union  of  chlorine  with 
the  metal  (cf.  p.  74),  or  by  the  combined  action  of  carbon  and  chlorine 
upon  the  oxide  (cf.  p.  520).  The  latter  method  is  used  in  making 
chromium  chloride.  The  general  methods  for  making  any  salt  (p.  186), 
such  as  the  interaction  of  a  metal  with  an  acid,  or  of  the  oxide,  hydrox- 
ide, or  another  salt  with  an  acid,  or  the  double  decomposition  of  two 
salts,  may  be  used  also  for  making  chlorides.  The  chlorides  are  for 
the  most  part  soluble  in  water.  Silver  chloride,  mercurous  chloride, 
and  cuprous  chloride  are  almost  insoluble,  however,  and  lead  chloride 
is  not  very  soluble.  Most  of  the  chlorides  of  metals  dissolve  without 
decomposition,  but  hydrolysis  is  conspicuous  in  the  case  of  the  chlorides 
of  the  trivalent  metals,  such  as  aluminium  chloride  and  ferric  chlo- 
ride (cf.  p.  344).  The  chlorides  of  some  of  the  bivalent  metals  are 
hydrolyzed  also,  but,  as  a  rule,  only  when  they  are  heated  with  water. 
This  is  the  case  with  the  chlorides  of  magnesium,  calcium,  and  zinc. 
Most  of  the  chlorides  are  stable  when  heated,  but  those  of  the  noble 
metals,  particularly  gold  and  platinum,  are  decomposed,  and  chlorine 
escapes.  The  chlorides  are  usually  the  most  volatile  of  the  salts  of  a 
given  metal,  and  so  are  preferred  for  the  production  of  the  spectrum 
(q.v.)  of  the  metal,  and  for  fixing  the  atomic  weight  of  the  metal  by 
use  of  the  vapor  density.  Some  of  the  metals  form  two  or  more  differ- 
ent chlorides.  For  example,  indium  gives  InCl,  InCl2,  and  InCl3. 

The  sulphides  are  formed  by  the  direct  union  of  the  metal  with  sul- 
phur, or  by  the  action  of  hydr  ogen  sulphide  or  of  some  soluble  sulphide 
upon  a  solution  of  a  salt  (cf.  p.  375).  In  one  or  two  cases  they  are  made 
by  the  reduction  of  the  sulphate  with  carbon.  The  sulphides,  except 
those  of  the  alkali  metals,  are  but  little  soluble  in  water.  The  sul- 
phides of  aluminium  and  chromium  are  hydrolyzed  completely  by 
water,  giving  the  hydroxides,  and  those  of  the  metals  of  the  alkaline 
earths  are  partially  hydrolyzed  (cf.  p.  376). 

Some  of  the  metals,  when  they  are  in  the  molten  form,  simply  dis- 
solve carbon,  ind,  when  they  are  cooled  once  more,  deposit  it  in  the 


THE   BASE-FORMING  ELEMENTS  548 

form  of  graphite.  This  is  true  particularly  of  platinum  and  iron. 
The  carbides  are  usually  formed  in  the  electric  furnace  by  interaction 
of  an  oxide  with  carbon  (cf.  p.  478).  Some  of  them  are  decomposed  by 
contact  with  water,  after  the  manner  of  calcium  carbide,  giving  a 
hydroxide  and  a  hydrocarbon.  Of  this  class  are  lithium  carbide  Li2C2, 
barium  and  strontium  carbides  BaC2  and  SrC2,  aluminium  carbide 
A14C3,  manganese  carbide  MnC,  and  the  carbides  of  potassium  and 
glucinum.  Others,  such  as  those  of  molybdenum  CMo2  and  chromium 
CrsC2,  are  not  affected  by  water. 

The  nitrates  may  be  made  by  any  of  the  methods  used  for  preparing 
salts.  They  are  all  at  least  fairly  soluble  in  water. 

The  sulphates  are  made  by  the  methods  used  for  making  salts,  and 
in  some  cases  by  the  oxidation  of  sulphides.  They  are  all  soluble 
in  water,  with  the  exception  of  those  of  lead,  barium,  and  strontium. 
Calcium  sulphate  is  meagerly  soluble. 

The  carbonates  are  prepared  by  the  methods  used  for  making  salts. 
They  are  all  insoluble  in  water,  with  the  exception  of  those  of  sodium 
and  potassium.  The  hydroxides  of  aluminium  and  tin  are  so  feebly 
basic  that  these  metals  do  not  form  stable  carbonates  (c/.pp.  119,  523). 

The  phosphates  and  silicates  are  prepared  by  the  methods  used  in 
making  salts.  The  former  are  obtained  also  by  special  processes 
already  described  (p.  467).  With  the  exception  of  the  salts  of  sodium 
and  potassium,  all  the  salts  of  both  these  classes  are  insoluble. 

Solubility  of  Bases  and  Salts.  —  The  solubilities  of  a  few  salts 
at  various  temperatures  have  already  been  given  (p.  157).  The  fol- 
lowing table  includes  a  much  larger  number  of  substances  (142),  and 
gives  the  solubility  at  18°.  Each  square  contains  two  numbers  ex- 
pressing the  solubility  of  the  compound  whose  cation  stands  at  the 
head  of  the  column  and  whose  anion  is  indicated  at  the  side.  The 
solubility  is  that  of  the  hydrate  stable  at  18°,  where  such  exists. 
The  upper  number  in  each  case  gives  the  number  of  grams  of  the  an- 
hydrous salt  held  in  solution  by  100  c.c.  of  water.  The  lower  number 
shows  the  number  of  moles  in  1  1.  of  the  saturated  solution,  and  indi- 
cates therefore  the  concentration  in  terms  of  a  molar  solution  as 
unity  —  the  molar  solubility.  In  the  cases  of  the  less  soluble  com- 
pounds the  values  are  not  exact,  but  they  will  serve  to  show  roughly 
the  relative  solubilities  when  several  substances  are  compared.  The 
numbers  for  small  solubilities  have  been  abbreviated.  Thus,  0.064  = 
0.0000004. 


544  INORGANIC   CHEMISTRY 

Solubility  of  Bases  and  Salts  in  Water  at  18°. 


K 

Na 

Li 

Ag 

Tl 

Ba 

Sr 

Ca 

Mg 

Zn 

Pb 

Cl 

32.95 
3.9 

35.86 
5.42 

77.79 
13.3 

0.0316 
0.0410 

0.3 
0.013 

37.24 
1.7 

51.09 
3.0 

73.19 
5.4 

55.81 
5.1 

203.9 
9.2 

1.49 
0.05 

Br 

65.86 
4.6 

88.76 
6.9 

168.7 
12.6 

0.041 
0.066 

0.04 
0.0215 

103.6 
2.9 

96.52 
3.4 

143.3 
5.2 

103.1 
4.6 

478.2 
9.8 

0.598 
0.02 

I 

137.5 
6.0 

177.9 

8.1 

161.5 
8.5 

0.063" 
0.071 

0.006 
0.0317 

201.4 
3.8 

169.2 
3.9 

200 
4.8 

148.2 
4.1 

419 
6.9 

0.08 
0.022 

F 

92.56 
12.4 

4.44 
1.06 

0.27 
0.11 

195.4 
13.5 

72.05 
3 

0.16 
0.0292 

0.012 
0.001 

00016 
0.032 

0.0076 
0.0214 

0.005 
0.0S5 

0.07 
0.003 

NO, 

30.34 
2.6 

83.97 
7.4 

71.43 
7.3 

213.4 

8.4 

8.91 
0.35 

8.74 
0.33 

66.27 
2.7 

121.8 
5.2 

74.31 
4.0 

117.8 

4.7 

51.66 
1.4 

C1OS 

6.6 
0.52 

97.16 
6.4 

313.4 
15.3 

12.25 
0.6 

3.69 
0.13 

35.42 
1.1 

174.9 
4.6 

179.3 
5.3 

126.4 

4.7 

183.9 
5.3 

150.6 
3.16 

Br03 

6.38 
0.38 

36.67 
2.2 

152.5 
8.20 

0.59 
0.025 

0.30 
0.009 

0.8 
0.02 

30.0 
0.9 

85.17 
2.3 

42.86 
1.5 

58.43 
1.8 

1.3 
0.03 

ios 

7.62 
0.35 

8.33 
0.4 

80.43 
3.84 

0.004 
0.0314 

0.059 
0.0216 

0.05 
0.001 

0.25 
0.0257 

0.25 
0.007 

6.87 
0.26 

0.83 
0.02 

0.002 
0.043 

OH 

142.9 
18 

116.4 
21. 

12.04 
5.0 

0.01 
0.001 

40.04 
1.76 

3.7 
0.22 

0.77 
0.063 

0.17 
0.02 

0.001 
0.032 

0.035 
0.045 

0.01 
0.034 

S04 

11.11 
0.62 

16.83 
1.15 

35.64 

2.8 

0.55 
0.020 

4.74 
0.09 

0.0323 
0.0410 

0.011 
0.036 

0.20 
0.015 

35.43 

2.8 

53.12 
3.1 

0.0041 
0.0313 

Cr04 

63.1 

2.7 

61.21 
3.30 

111.6 
6.5 

0.0025 
0.0315 

0.006 
0.031 

0.0338 
0.0415 

0.12 
0.006 

0.4 
0.03 

73.0 
4.3 

0.042 
0.0«5 

C204 

30.27 
1.6 

3.34 
0.24 

7.22 
0.69 

0.0035 
0.032 

1.48 
0.030 

0.0086 
0.0338 

0.0046 
0.0326 

0.0356 
0.0443 

0.03 
0.0027 

0.036 
0.044 

0.0315 
0.0r,5 

cos 

108.0 
5.9 

19.39 
1.8 

1.3 
0.17 

0.003 
0.031 

4.95 
0.10 

0.0023 
0.0311 

0.0011 
0.047 

0.0013 
0.0313 

0.1 
0.01 

0.004? 
0.033? 

0.031 
0.043 

The  following  are  the  solubilities  (number  of  grams  in  100  c.c.  water 
at  18°  of  two  additional  insoluble  substances  and  of  three  acid  salts 


Mercurous  chloride,  0.032 
(molar  soPty,  0.041) 
Mercuric  iodide,       0.044 
Qmoiar  sol'ty,  0.051) 


Sodium  bicarbonate  9.6 

Potassium  bicarbonate     26.1 
Potassium  bisulphate       50.0 


THE   BASE-FORMING  ELEMENTS 


545 


It  will  be  seen  that  some  compounds,  like  zinc  chloride  and  barium 
iodide,  are  exceedingly  soluble ;  that  others,  like  potassium  chloride  and 
barium  chlorate,  are  of  medium  solubility  ;  that  still  others,  like  calcium 
hydroxide  and  calcium  sulphate,  are  sparingly  soluble  ;  and,  finally,  that 
some,  like  calcium  oxalate  (CaC2O4)  and  barium  chromate,  are  almost 
insoluble.  The  reader  should  note  the  fact,  however,  that  the  differ- 
ences in  solubility  even  amongst  the  insoluble  salts  are  as  great  as 
amongst  the  soluble  ones. 

Hydrated  Forms  of  Salts  Commonly  Used.  —  In  the  table 
given  below,  the  figures  refer  to  the  number  of  molecules  of  water  in 
the  hydrates  which  are  deposited  by  aqueous  solutions  of  the  salts  in 
the  neighborhood  of  18°.  The  letter  h  means  that  the  compound  is 
stable  when  heated,  the  letter  a  that  it  is  not  affected  by  the  air,  the 
letter  d  that  the  salt  is  deliquescent,  and  the  letter  e  that  the  hydrate 
loses  water  spontaneously  in  an  open  vessel,  i.e.,  is  efflorescent. 

Composition  of  Hydrates  of  Salts. 


K 

Na 

Li 

Ag 

Ba 

Sr 

Ca 

Mg 

Zn 

Cd 

Cu 

Pb 

Cl  . 

Oh 

Oh 

Oh 

Oh 

2a 

Qe 

6d 

6d 

lid 

2^e 

2d 

Oh 

Br  . 

Oh 

Oh 

Oh 

Oh 

2a 

Qd 

6d 

6d 

2d 

4e 

4e 

Oh 

I   . 

Oh 

Oh 

Oh 

Oh 

2d 

Qd 

Od 

8d 

Od 

Oa 

. 

Oh 

NO, 

Oa 

Oa 

Od 

Oa 

Oa 

4e 

4d 

Qd 

Qd 

4d 

6d 

Oa 

C103 

Oa 

Oa 

Od 

Oa 

la 

5d 

2d 

Qd 

Qd 

2d 

Qd 

la 

BrO3 

Oa 

Oa 

Od 

Oa 

la 

la 

la 

Qe 

Qa 

2a 

Qa 

la 

IO3 

Oa 

3e 

Od 

Oa 

la 

Qe 

Oe 

4a 

2a 

Oa 

la 

Oa 

C2H302 

Od 

3d 

2d 

Oa 

la 

$a 

2d 

4d 

3a 

3d 

la 

3a 

S04. 

Oh 

lOe 

Oh 

Oa 

Oh 

Oh 

2a 

7e 

7e 

2|a 

5a 

Oh 

Cr04 

Oh 

lOe 

2a 

Oa 

Oh 

Oh 

la 

7e 

3 

Oh 

C204 

la 

Oa 

Oa 

Oa 

la 

Oa 

la 

2a 

2a 

3a' 

la 

Oa 

C03 

lid 

lOe 

Oa 

Oe 

Oh 

Oh 

Oa 

3e 

Oa 

Oa 

Oa 

Isomorphism.  —  Substances  which  crystallize  in  one  of  the  forms 
belonging  to  the  regular  syStem  (p.  138)  must  necessarily  have  identi- 
cal crystalline  shapes.  Thus,  crystallized  specimens  of  sodium  chlo- 
ride and  of  lead  sulphide  (galena)  in  their  natural  shapes  are  cubical. 
The  forms  found  in  other  systems,  however,  are  capable  of  assuming 
an  infinite  diversity  of  shapes.  The  relative  lengthening  or  shorten- 
ing in  one  direction,  shown  by  the  square  prismatic  and  hexag- 
onal forms  (p.  138),  for  example,  makes  it  possible  for  each  separate 
substance  to  adopt  proportions  which  are  more  or  less  different 


546  INORGANIC   CHEMISTRY 

from  those  of  every  other  substance.  Each  substance,  not  belonging 
to  the  regular  system,  does,  in  fact,  crystallize  invariably  in  forms 
based  upon  its  own  fundamental  proportions,  and  differs  therefore  in 
its  angles  from  all  other  such  substances  in  a  way  that  is  clearly 
recognizable  by  refined  measurement. 

Now  it  is  found  that  substances  which  are  chemically  somewhat 
similar  (see  below)  frequently  crystallize  in  the  same  system  and  show 
proportions  which  are  almost,  although  not  quite,  identical.  Further- 
more, such  substances,  when  the  approach  to  identity  in  angles  is  not 
accidental,  can  take  part  in  the  construction  of  one  and  the  same 
crystal.  A  crystal  of  one  such  substance  placed  in  a  solution  of  the 
other  will  continue  to  grow,  and  in  doing  so  will  follow  the  pattern 
already  set,  and  simply  increase  in  dimensions  by  accretion  of  the  new 
material.  When  a  solution  containing  two  such  substances  deposits 
crystals,  the  structures  are,  not  some  of  them  of  one  material  and  some 
of  the  other,  but  are  all  made  up  of  both  in  a  ratio  determined  by  the 
relative  amounts  of  the  substances  in  the  solution.*  Substances 
related  in  these  two  ways,  that  is,  baring,  when  separate,  crystalline 
forma  -which  are  closely  alike  and  being  capable  of  forming  homoge- 
neous crystals  containing  varying  proportions  of  the  two  ingredients, 
are  called  isomorphous  substances  (Gk.  to-os,  equal;  /no/Deform). 
Thus,  potassium  permanganate  KMn04  and  potassium  perclilorate 
KC104  crystallize  in  the  rhombic  system,  forming  crystals  with  very 
similar  angles  (Fig.  53,  p,  139),  and  when  a  solution  containing  both  is 
allowed  to  evaporate  there  is  formed  but  one  set  of  crystals  made  up  of 
both  substances.  Similarly,  potassium  iodide  and  ammonium  iodide 
crystallize  in  cubes  of  the  regular  system,  and,  since  all  cubes  are 
alike,  necessarily  show  absolutely  identical  angles.  In  addition  to  this, 
however,  they  crystallize  together  from  a  solution  containing  both 
salts.  Other  pairs  from  amongst  substances  belonging  to  the  regular 
system  would  not  do  this.  The  two  salts  are  therefore  isomorphous. 

In  the  course  of  our  study  of  the  compounds  of  the  metals  we 
shall  have  occasion  to  note  many  examples 'of  isomorphism.  Thus  the 
heptahydrates  of  the  sulphates  of  many  of  the  bivalent  metals,  such 
as  ZnS04,7H2O,  FeS04,7H2O,  MgS04,7H2O,  etc.,  belong  to  the  rhombic 
system,  and  form  an  isomorphous  set  of  substances  known  as  the 
vitriols  (q.v.). 

*  In  general,  two  substances  which  are  .absolutely  unrelated  may  be  deposited 
simultaneously  from  mixed  aqueous  solutions,  but  some  of  the  crystals  are  pure 
specimens  of  one  substance,  and  the  rest  are  pure  specimens  of  the  other. 


THE   BASE-FORMING  ELEMENTS  547 

The  alums  (£.#.)  also  constitute  an  important  set,  and  crystallize, 
separately  and  together,  in  the  regular  system  (see,  also,  under  Zinc 
sulphate).  Amongst  minerals,  lead  sulphide  (PbS)  and  silver  sulphide 
(Ag2S)  form  a  common  isomorphous  pair,  and  nearly  all  natural  speci- 
mens of  galena  (q.v.)  contain  at  least  a  little  silver  sulphide. 

The  chemical  significance  of  isomorphism  was  at  first  exaggerated. 
Thus  the  elements  magnesium  and  iron  are  not  especially  similar  in 
their  chemical  relations,  excepting  that  both  are  bivalent ;  yet  they 
form  several  pairs  of  compounds  which,  like  the  sulphates  (above),  are 
isomorphous.  Still,  in  practical  chemical  work  a  knowledge  of  the 
relations  of  a  substance  in  respect  to  isomorphism  is  indispensable. 
It  enables  us  to  predict  the  probable  impurities  in  a  homogeneous- 
looking  material,  for  non-isomorphous  substances  would  have  given  a 
heterogeneous  mixture  with  it.  It  assists  us  in  separating  and  puri- 
fying chemical  substances,  for  non-isomorphous  substances  can  be  sep- 
arated by  recrystallization  from  water  (p.  273),  or  by  washing  with 
water  or  some  other  solvent  (p.  276),  at  a  temperature  at  which  the 
solubilities  of  the  substances  are  different  (see  Potassium  nitrate). 
Isomorphous  substances,  however,  can  be  separated  only  by  conversion 
into  some  other  form  of  combination  in  which  the  property  is  lack- 
ing. Thus,  silver  sulphide  cannot  be  separated  from  lead  sulphide  by 
Pattinson's  process  (q.v.),  and  so  the  mixed  metals,  to  the  separation 
of  which  the  process  is  applicable,  must  first  be  secured  by  reduction. 

Some  chemists  regard  isomorphous  mixtures  as  solid  solutions. 

Exercises.  —  1.  Compare  the  electrical  conductivities  of  normal 
sodium  hydroxide  and  normal  acetic  acid  with  the  conductivity  of 
copper.  What  length  of  copper  wire  will  present  the  same  resistance 
as  1  cm.  of  each  of  these  solutions  when  the  cross-sections  are  alike  ? 

2.  What  do  we  mean  by  saying  that  an  oxide  is  strongly  or  feebly 
basic,  or  that  it  is  acidic  (p.  541)  ? 

3.  What  is  meant  by  the  same  terms  when  applied  to  an  hydroxide  ? 

4.  Compare  the  molar  solubilities  at  18°  (a)  of  the  halides  of  silver 
and  (b)  of  the  carbonates  and  (c)  oxalates  of  the  metals  of  the  alkaline 
earths,  noting  the  relation  between  solubility  and  atomic  weight. 

5.  What  is  the  molar  concentration  of  chloridion  (cf.  p.  149)  in 
saturated  solutions  of  silver  chloride  and  lead  chloride  at  18°,  assum- 
ing complete  ionization  in  these  very  dilute  solutions  ? 

6.  How  does  the  behavior  of  complex  acids,  like   chloroplatinio 
acid  HaPtCl6,  differ  from  that  of  acid  salts  ? 


CHAPTER  XXXTII 


THE   METALS    OF   THE    ALKALIES:    POTASSIUM    AND 
AMMONIUM 

The  Metals  of  the  Alkalies.  —  The  metals  of  this  family  form  a 
homogeneous  group,  and  there  is  a  very  general  similarity  between  the 
properties  of  the  corresponding  compounds.  Some  of  the  physical 
properties  of  the  elements  themselves  can  be  presented  best  in  tabular 
form. 


AT.  WT. 

SP.  GB. 

M.-P. 

B.-P. 

Lithium  . 

7.0 

0.53 

186° 

above  red  heat. 

Sodium  . 

23.0 

0.97 

95.6° 

742° 

Potassium 

39.1 

0.86 

62.5° 

667° 

Rubidium 

85.5 

1.53 

38.5° 

Caesium  . 

133.8 

1.87 

26.5° 

270° 

It  will  be  seen  that  the  specific  gravities  of  the  elements  increase  with 
rising  atomic  weight,  while  the  melting-points  and  boiling-points  fall 
(cf.  p.  410).  A  table  including  all  the  physical  properties,  both  of  the 
elements  and  their  compounds,  would  show  similar  characteristics  in  a 
general  way,  with  here  and  there  noticeable  irregularities  such  as  that 
shown  by  the  specific  gravity  of  sodium. 

The  Chemical  Relations  of  the  Metallic  Elements  of  the  Alka- 
lies. —  The  metals  which  are  chemically  most  active  are  included  in 
this  group,  and  the  activity  increases  with  rising  atomic  weight,  caesium 
being  the  most  active  positive  element  of  all.  A  freshly  cut  surface  of 
any  of  these  metals  tarnishes  by  oxidation  as  soon  as  it  is  exposed  to  the 
air.  Indeed,  there  is  scarcely  time  to  see  the  metallic  appearance  in 
the  case  of  potassium  and  the  metals  following  it.  All  of  these  metals 
decompose  water  violently  (cf.  p.  97),  liberating  hydrogen.  The 
hydroxides  which  are  formed  by  this  action  are  exceedingly  active 
bases,  that  is  to  say,  they  give  a  relatively  large  concentration  of 
Irydroxidion  in  solutions  of  a  given  molecular  concentration  (p.  349). 

548 


POTASSIUM  AND  AMMONIUM  549 

Lithium  hydroxide  is  the  least  active.  In  the  dry  form  these  hydrox- 
ides are  not  decomposed  by  heating,  while  the  hydroxides  of  all  other 
metals  lose  water  more  or  less  easily.  All  these  metals  seem  to  com- 
bine with  hydrogen,  lithium  giving  the  most  stable  compound.  The 
hydrides,  however,  unlike  those  of  many  of  the  non-metals,  are  not 
ionogens,  and  consequently  do  not  give  acids  when  dissolved  in  water. 
In  all  their  compounds  the  metals  of  the  alkalies  are  univalent. 

The  family  may  be  subdivided  into  two  minor  groups.  The  com- 
pounds of  potassium,  rubidium,  and  caesium  resemble  one  another 
closely,  while  those  of  sodium  and  lithium  are  sometimes  largely  diver- 
gent in  physical  properties.  Thus,  the  chlorides  of  the  potassium  set,  not 
only  crystallize  in  cubes,  but  can  form  mixed  crystals  with  one  another 
in  all  proportions.  They  are  isomorphous  (cf.  p.  545).  The  same  is 
true  of  the  bromides  and  of  the  iodides.  Sodium  chloride,  although 
crystallizing  in  cubes  likewise,  does  not  form  mixed  crystals  with  the 
chlorides  of  the  potassium  set.  In  the  case  of  lithium,  the  hydroxide 
is  not  nearly  so  soluble  as  are  the  hydroxides  of  the  other  metals,  and 
the  metal  gives  also  an  insoluble  carbonate  and  phosphate,  in  which 
respect  it  resembles  magnesium  and  differs  from  all  the  other  mem- 
bers of  the  present  group. 

The  compounds  of  ammonium  will  be  discussed  in  connection  with 
those  of  potassium,  to  which  they  present  the  greatest  resemblance. 

The  solubilities  are  often  decisive  factors  in  connection  with  the 
preparation  and  use  of  salts.  The  reader  will  find  most  of  these  in  the 
table  on  p.  544,  or  the  diagram  on  p.  157,  and,  as  a  rule,  the  values  will 
not  be  repeated  in  the  descriptive  paragraphs. 

POTASSIUM. 

Occurrence.  —  Silicates  containing  potassium,  such  as  feldspar  and 
mica  (p.  525),  are  constant  constituents  of  volcanic  rocks,  and  from  the 
weathering  of  these  rocks,  and  of  the  detritus  formed  from  them  which 
constitutes  a  large  part  of  the  soil,  the  potassium  used  by  plants  is 
obtained.  These  minerals  are  not  used  commercially  as  sources  of 
potassium  compounds.  The  salt  deposits  (see  below)  contain  potassium 
chloride,  alone  (sylvite)  and  in  combination  with  other  salts,  and  most 
of  the  compounds  of  potassium  are  manufactured  from  this  material. 
Part  of  our  potassium  nitrate,  however,  is  purified  Bengal  saltpeter 
(p.  438).  Potassium  sulphate  occurs  also  in  the  salt  layers,  and  is  used 
directly  as  a  fertilizer. 


550  INORGANIC   CHEMISTRY 

Preparation.  —  Potassium  was  first  made  by  Davy  (1807)  by 
bringing  the  wires  from  a  battery  in  contact  with  a  piece  of  moist 
potassium  hydroxide.  Globules  of  the  metal  appeared  at  the  negative 
wire.  This  process  has  just  come  into  use  commercially,  molten  potas- 
sium chloride  being  the  substance  decomposed.  The  process  which, 
during  the  intervening  years,  furnished  potassium,  was  the  heating  of 
potassium  carbonate  with  finely  divided  carbon  in  small  retorts  : 

K2C03  +  2C  ->  2K  +  SCO. 

The  vapor  of  the  metal  tends  to  combine  with  the  carbon  monoxide, 
forming  an  explosive  compound  K6C6O6,  and  the  yield  is  thus  reduced. 
Castner's  improved  process  involves  the  heating  of  potassium  hydroxide 
with  a  spongy  mass  which  is  essentially  a  carbide  of  iron  (CFe2).  The 
latter  is  made  by  heating  together  pitch  and  iron  filings.  No  carbon 
monoxide  is  produced : 

6KOH  +  20  ->  2K2C03  +  3H2  +  2K. 

i 
Physical  Properties.  —  Potassium  is  a  silver- white  metal  which 

melts  at  62.5°.  It  boils  at  667°,  giving  a  greenish  vapor.  The  metal 
and  its  compounds  confer  a  violet  tint  upon  the  Bunsen  flame,  and  the 
spectrum  (q.v.)  shows  characteristic  lines. 

Chemical  Properties.  —  The  density  of  the  vapor  shows  the 
molecular  weight  of  potassium  to  be  about  40,  so  that  the  vapor  is  a  mon- 
atomic  gas.  The  element  unites  violently  with  the  halogens,  sulphur, 
and  oxygen.  In  consequence  of  the  latter  fact  it  is  usually  kept  under 
petroleum,  an  oil  which  neither  contains  oxygen  itself,  nor  dissolves  a 
sufficient  amount  of  oxygen  from  the  air  to  permit  much  oxidation  of 
the  potassium  to  take  place. 

The  Hydride.  —  When  hydrogen  is  passed  over  potassium  heated 
to  360°,  a  hydride  is  formed.  By  washing  the  solid  product  with 
liquefied,  dry  ammonia  the  excess  of  potassium  is  removed  and  white 
crystals  remain.  These  have  the  composition  KH.  On  account  of  the 
ease  with  which  it  decomposes,  the  substance  behaves  much  like 
potassium  itself.  When  thrown  into  water,  for  example,  it  gives 
potassium  hydroxide,  and  the  hydrogen  is  liberated. 

Potassium  Chloride.  —  Sea-water  and  the  waters  of  salt  lakes 
contain  a  relatively  small  proportion  of  potassium  compounds.  During 


POTASSIUM  AND  AMMONIUM  551 

the  evaporation  of  such  waters,  however,  the  potassium  compounds 
tend  to  accumulate  in  the  mother-liquor  while  sodium  chloride  is  being 
deposited.  Thus,  when  the  Salt  Lake  in  Utah  shall  have  finally  dried 
up,  the  upper  (last  to  be  formed)  part  of  the  bed  of  salts  which  it  will 
leave  behind  will  contain  layers  rich  in  compounds  of  potassium.  This 
condition  is  realized  in  geological  deposits  which  have  been  formed  in 
the  same  way.  Thus,  at  Stassfurt,  near  Magdeburg,  there  is  a  thickness 
of  more  than  a  thousand  meters  of  common  salt,  more  or  less  mixed 
with  and  intersected  by  layers  of  sedimentary  deposits.  Above  this  are 
25 -30  meters  of  salt  layers  in  which  the  potassium  salts  are  chiefly 
found,  while  over  all  are  several  hundred  meters  of  sandstone.  For- 
merly the  upper  layers  were  simply  stripped  off  and  rejected.  Now, 
however,  the  revenue  obtained  from  the  products  of  these  layers  is 
many  times  greater  than  that  coming  from  the  rock  salt  below. 

The  chief  sources  of  the  potassium  chloride  found  in  the  salt  beds 
are  sylvite  (KC1)  and  carnallite  (KC1,  MgCl2,  6H20).  The  latter  is 
heated  with  a  small  amount  of  water,  or  with  a  mother-liquor  obtained 
from  a  previous  operation  and  containing  sodium  and  magnesium 
chlorides.  The  magnesium  sulphate  which  it  contains  as  an  impurity 
remains  undissolved.  From  the  clear  liquid,  when  it  cools,  potassium 
chloride  is  deposited  first  and  then  carnallite.  The  former  is  taken  out 
and  purified,  and  the  latter  goes  through  the  process  again.  This 
potassium  chloride  is  the  source  from  which  mott  of  our  potassium 
hydroxide  and  potassium  carbonate,  as  well  as  salts  of  minor  commer- 
cial importance,  are  -made.  It  is  a  white  substance  crystallizing  in 
cubes,  melting  at  about  750°,  and  slightly  volatile  at  high  temperatures. 

Potassium  Iodide.  —  When  iodine  is  heated  in  a  strong  solution 
of  potassium  hydroxide,  the  iodate  and  iodide  are  both  formed  (p.  277) : 

6KOH  +  3I2  ->  5KI  +  KIO3  +  3H2O. 

The  dry  residue  from  evaporation  is  heated  with  powdered  carbon  to 
reduce  the  iodate,  and  all  the  iodide  can  then  be  purified  by  recrystal- 
lization.  Another  method  of  preparation  consists  in  rubbing  together 
iodine  and  iron  filings  under  water.  The  soluble  ferrous  iodide  (FeI2) 
thus  formed  is  then  treated  with  additional  iodine  and  gives  a  sub- 
stance Fe3I8,  intermediate  in  composition  between  ferrous  and  ferric 
iodides.  This  is  also  soluble.  When  potassium  carbonate  is  added  to 
the  solution,  a  hydrated  magnetic  oxide  of  iron  is  precipitated,  carbon 


552  INORGANIC   CHEMISTRY 

dioxide  escapes,  and  evaporation  of  the  clear  solution  gives  potassium 
iodide : 

Fe8I8  +  4K2CO3  +  4HaO  -»  SKI  +  Fe8(OH)8  +  4C02. 

The  salt  forms  large,  somewhat  opaque  cubes  (rn.-p.  623°).  It  is 
used  in  medicine  and  for  producing  precipitates  of  silver  iodide  in 
photography.  In  the  laboratory  it  is  used  whenever  an  iodide  is 
required,  for  example,  when  experiments  with  iodidion  are  to  be  made 

The  Bromide  and  Fluorides.  —  Potassium  bromide  may  be 
made  in  either  of  the  ways  used  for  the  iodide.  It  crystallizes  in 
cubes.  It  is  used  in  medicine  and  for  precipitating  silver  bromide  in 
making  photographic  plates  (q.v.).  In  the  laboratory  it  is  always 
employed  when  a  bromide  is  needed  as  a  source  of  bromidion. 

The  fluoride  of  potassium  K2F2  may  be  obtained  by  treating  the 
carbonate  or  hydroxide  with  hydrofluoric  acid.  It  is  a  deliquescent, 
white  salt.  When  treated  with  an  equi-molecular  quantity  of  hydro- 
fluoric acid  it  forms  potassium  hydrogen  fluoride  KHF2,  a  white  salt 
which  is  also  very  soluble.  This  acid  salt  is  used  in  the  preparation 
of  pure  hydrofluoric  acid,  since  the  latter  is  liberated  from  it  as  a  vapor 
at  a  high  temperature. 

Potassium  chloride  is  the  least  soluble  of  the  halides  of  potassium, 
the  bromide,  fluoridt,  and  iodide  coining  next  in  that  order.  The  posi- 
tion of  the  fluoride  as  the  third  in  order,  when  we  should  expect  it  to  be 
the  least  soluble  (p.  244),  shows  that  this  compound  is  somewhat 
exceptional.  It  is  also  slightly  hydrolyzed  by  water,  as  if  it  were  a 
salt  of  a  dibasic  acid  (cf.  p.  344).  These  facts,  together  with  the 
existence  of  the  acid  fluoride,  lead  us  to  assign  to  it*  the  formula  K2F2. 
Other  acid  fluorides  of  the  formulae  KH2F8  and  KH8F4  have  likewise 
been  made.  Since  potassium,  hydrogen,  and  fluorine  are  always  univa- 
lent,  and  no  ordinary  valence  is  thus  available  for  holding  together 
groupings  more  complex  than  KF  and  HF,  we  may  regard  all  these 
four  fluorides  of  potassium  as  molecular  compounds  (p.  443). 

Potassium  Hydroxide*  —  This  compound,  known  also  as  caustic 
potash  and  sometimes  as  potassium  hydrate  (p.  120),  was  formerly 
made  entirely  by  boiling  potassium  carbonate  with  calcium  hydroxide 
suspended  in  water  (milk  of  lime)  : 

K2C08  +  |  Ca(OH)2  <=±  CaC08  |  +  2KOH. 


POTASSIUM   AND   AMMONIUM  553 

The  operation  is  conducted  in  iron  vessels,  because  porcelain,  being 
composed  of  silicates,  interacts  with  solutions  of  bases.  The  action  is, 
in  theory,  precisely  similar  to  that  of  sulphuric  acid  upon  barium  diox- 
ide (cf.  p.  585).  The  potassium  carbonate  corresponds  to  the  acid, 
being  completely  dissolved  from  the  beginning,  and  the  calcium  hydrox- 
ide to  the  dioxide,  since  its  relative  insolubility  enables  the  water  to 
take  up  fresh  portions  into  solution  only  when  the  part  dissolved  has 
already  undergone  chemical  change.  The  calcium  carbonate  which  is 
precipitated  is  much  less  soluble  than  the  hydroxide,  and  hence  the 
action  goes  forward.  The  action  as  a  whole  is  reversible,  for  a 
reason  which  will  be  explained  later  (see  Ionic  equilibrium  in 
Chap,  xxxiv),  and  consequently  such  an  amount  of  water  is  employed 
that  the  solution  at  no  time  contains  more  than  about  ten  per  cent  of 
potassium  hydroxide  (sp.  gr.  1.1).  The  conclusion  of  the  action  is 
recognized  when  a  clear  sample  of  the  liquid  no  longer  effervesces 
on  addition  of  a  dilute  acid,  and  is  therefore  free  from  potassium 
carbonate. 

Recently  much  potassium  hydroxide  has  been  manufactured  by 
electrolytic  processes.  When  a  solution  of  potassium  chloride  is  elec- 
trolyzed,  chlorine  is  liberated  at  the  anode,  and  hydrogen  and  potassium 
hydroxide  at  the  cathode.  The  necessity  of  keeping  those  two  sets  of 
products  apart,  since  by  their  interaction  potassium  hypochlorite  and 
potassium  chloride  would  be  formed  (cf.  p.  266),  has  made  the  devising 
of  suitable  apparatus  extremely  difficult.  In  one  type  of  apparatus  a 
partition  of  asbestos  cloth,  especially  prepared  to  resist  the  disinte- 
grating effects  of  the  alkali  and  the  chlorine,  divides  the  cell  into  two 
parts.  In  some  cases  this  is  placed  vertically,  and  in  others  horizon- 
tally. In  the  latter  case  the  anode  is  on  the  upper  side  of  the  partition, 
in  order  that  the  chlorine  as  it  is  liberated  may  ascend  to  the  surface 
without  stirring  up  the  liquid  or  having  occasion  to  pass  near  the  par- 
tition. In  all  cases  the  anode  is  made  of  graphite,  since  this  substance 
is  less  easily  attacked  by  chlorine  than  is  any  other,  and  the  cathode  is 
made  of  iron,  a  metal  which  best  resists  the  action  of  alkalies.  The 
chlorine  is  used  for  making  bleaching  powder.  Pure  brine  flows  in  con- 
tinuously at  one  point,  and  a  solution  of  the  hydroxide  containing  much 
undecomposed  chloride  flows  out  at  another. 

The  same  process  is  applied  to  sodium  chloride,  and,  in  some  fac- 
tories, the  apparatus  is  used  solely  for  making  bleaching  materials,  and 
most  or  all  of  the  alkali  is  thrown  away.  This  is  on  account  of  the 
cheapness  of  the  caustic  soda  made  by  non-electrolytic  processes,  and 


554 


INORGANIC  CHEMISTRY 


the  expense  involved  in  concentrating  the  rather   dilute    solution 
obtained  electrolytically. 

The  Castner-Kellner  apparatus  (Fig.  97)  employs  a  different  princi- 
ple very  ingeniously  for  the  separation  of  the  products.  The  two  end 
compartments  are  filled  with  brine  and  contain  the  graphite  anodes. 
The  central  compartment  contains  potassium  hydroxide  solution  and 
the  iron  cathode.  The  positive  current  enters  by  the  anodes,  and  the 
chlorine  is  therefore  attracted  to  and  liberated  upon  the  graphite. 
After  rising  through  the  liquid  it  is  collected  for  the  manufacture 
of  liquefied  chlorine  or  of  bleaching  powder.  The  ions  of  potassium 
or  of  sodium,  as  the  case  may  be,  are  discharged  upon  a  layer  of  mer- 


FIQ.  97. 


cury  which  covers  the  whole  floor  of  the  box,  and  the  free  metal 
dissolves  in  the  mercury,  forming  an  amalgam  (p.  532).  The  layer  of 
mercury  extends  beneath  the  partitions,  and  a  slight  rocking  motion 
given  to  the  cell  by  the  cam  (C)  causes  the  amalgam  to  flow  below 
the  partition  into  the  central  compartment.  Here  the  sodium  leaves 
the  mercury  in  the  form  of  sodium  ions  and  is  attracted  by  the 
cathode.  Upon  this,  hydrogen  from  the  water  is  discharged,  and  the 
residual  hydroxidion,  together  with  the  metallions,  constitutes  potas- 
sium or  sodium  hydroxide.  A  slow  influx  of  salt  solution  at  one 
point  and  overflow  of  the  alkaline  solution  in  the  central  cell  at 
another,  is  maintained.  The  overflowing  liquid  contains  20  per  cent 
of  the  alkali.  Since  in  this  form  of  the  apparatus  there  is  no  unde- 
composed  chloride  present  in  the  part  of  the  solution  which  contains 
the  hydroxide,  simple  evaporation  to  dryness  furnishes  the  solid 
alkali. 

Potassium  hydroxide  is  exceedingly  soluble  in  water,  and  conse- 
quently, instead  of  being  crystallized  from  solution,  the  molten  residue 


POTASSIUM   AND   AMMONIUM  555 

from  evaporation  is  cast  in  sticks.  When,  for  chemical  purposes,  the 
hydroxide  is  required  free  from  potassium  carbonate  and  other  impu- 
rities, it  is  dissolved  in  alcohol,  in  which  the  other  substances  are  not 
soluble.  Evaporation  of  this  solution  gives  pure  caustic  potash.  The 
hydroxide,  in  consequence  of  the  very  low  vapor  tension  of  its  solution 
(cf.  p.  162),  is  highly  deliquescent.  It  also  absorbs  carbon  dioxide  from 
the  air,  giving  potassium  carbonate.  This  salt  is  itself  deliquescent, 
and  consequently  a  syrupy  solution  of  the  carbonate  is  the  final  result 
of  weathering.  Solutions  of  the  hydroxide  have  an  exceedingly  corro- 
sive action  upon  the  flesh,  decomposing  it  into  a  slimy  mass  by  hydro- 
lyzing  the  albuminous  and  other  substances.  In  solution,  the  base  is 
highly  ionized,  furnishing  a  high  concentration  of  hydroxidion.  Its 
aqueous  solution  is  therefore  used  with  salts  of  other  metals  for  precipi- 
tation of  less  soluble  bases.  Commercially  it  is  chiefly  employed  in  the 
making  of  soft  soap. 

The  Oxides.  —  The  simple  oxide  K20  may  be  made  by  heating 
potassium  nitrate  or  nitrite  with  potassium  in  a  vessel  from  which  air 
is  excluded  :  KN03  +  5K  — >  3K2O  -j-  K  It  interacts  violently  with 
water,  giving  the  hydroxide.  When  exposed  to  the  air  it  unites  spon- 
taneously with  oxygen,  and  K2O4  is  formed. 

When  the  metal  burns  in  oxygen,  K2O4,  a  yellow  solid  is  the  prod- 
uct. This  substance  interacts  violently  with  water,  giving  potassium 
hydroxide,  and  the  excess  of  oxygen  is  liberated.  With  perfectly  dry 
oxygen,  potassium  does  not  unite,  even  when  it  is  heated  strongly. 

Potassium  Chlorate.  —  The  preparation  of  this  salt  (KC103),  by 
interaction  of  potassium  chloride  with  calcium  chlorate,  has  already 
been  described  (p.  273).  It  is  also  made  by  electrolysis  of  potassium 
chloride  solution,  the  potassium  hydroxide  and  chlorine  which  are 
liberated  being  precisely  the  materials  required.  All  that  is  necessary 
is  to  use  a  warm,  concentrated  solution  and  to  provide  for  the  mixing 
of  the  materials  generated  at  the  electrodes.  The  salt  crystallizes  out 
when  the  solution  cools. 

.  Potassium  chlorate  crystallizes  in  monoclinic  plates.  It  melts  at 
about  351°,  and  at  a  temperature  slightly  above  this  the  visible  libera- 
tion of  oxygen  begins.  Since  heat  is  given  out  by  the  decomposition, 
the  action  may  be  almost  explosive  if  large  amounts  of  the  material 
are  employed.  This  decomposition  doubtless  takes  place  at  all  tem- 
peratures (p.  73),  but  below  the  melting-point  the  speed  is  so  slight 


556  INORGANIC   CHEMISTRY 

that  the  phenomenon  is  not  perceived.  On  account  of  the  ease  with 
which  its  oxygen  is  liberated,  the  salt  is  employed  in  making  fire- 
works and  as  a  component,  along  with  antimony  trisulphide,  of  the 
heads  of  Swedish  matches.  With  acids  it  is  used  as  an  oxidizing  agent 
on  account  of  the  chloric  acid  which  is  set  free  (p.  272).  It  is  also  em- 
ployed in  medicine. 

Potassium  perchlorate  KC104,  formed  by  the  heating  of  the  chlorate 
(p.  275),  gives  white  crystals  belonging  to  the  rhombic  system.  Com- 
pared with  the  chlorate,  on  account  of  the  greater  difficulty  in  liberat- 
ing its  oxygen  by  heat,  it  finds  little  practical  application. 

The  Bromate  and  lodate.  — These  are  the  most  familiar  salts  of 
their  respective  acids.  The  mode  of  their  preparation  has  already  been 
described  (p.  277).  Potassium  iodate  may  be  made  also  very  conven- 
iently by  melting  together  potassium  chlorate  and  potassium  iodide  at 
a  low  temperature.  The  iodate  is  much  less  soluble  than  the  chloride, 
and  the  mixture  may  be  separated  by  crystallization  from  water. 

Potassium  Nitrate.  —  The  formation  of  this  salt  in  nature  and 
its  mode  of  extraction  and  purification  have  already  been  described 
(p.  438).  This  source  of  supply  proved  insufficient,  for  the  first  time, 
during  the  Crimean  war  (1852-55),  and  a  method  of  manufacture  from 
Chili  saltpeter  (sodium  nitrate),  which  is  a  much  cheaper  substance, 
was  introduced.  Sodium  nitrate  and  potassium  chloride  are  heated 
with  very  little  water,  and  the  sodium  chloride  produced  by  the  action, 
which  is  a  reversible  one,  is  by  far  the  least  soluble  of  the  four  salts.  On 
the  other  hand,  at  this  temperature,  the  potassium  nitrate  is  by  far  the 
most  soluble.  Hence  the  hot  liquid  drained  from  the  crystals  contains 
the  required  salt,  and  most  of  the  sodium  chloride  is  in  the  form  of  a 
precipitate.  If  the  solubility  curve  of  potassium  nitrate  (p.  157)  is  ex- 
amined, it  will  be  seen  that  this  salt  is  but  slightly  soluble  in  cold  water, 
and  hence  most  of  it  is  deposited  when  the  solution  cools.  The  crys- 
tals are  mixed  with  little  sodium  chloride,  for,  as  the  curve  shows,  com- 
mon salt  is  little  less  soluble  at  10°  than  it  is  at  100°. 

Potassium  nitrate  gives  long  prisms  belonging  to  the  rhombic  sys- 
tem (Fig.  98).  It  melts  at  about  340°,  and  when  more  strongly  heated 
gives  off  oxygen,  leaving  potassium  nitrite  (p.  449).  Although  it  does 
not  form  a  hydrate,  the  crystals  inclose  small  portions  of  the  mother- 
liquor,  and  consequently  contain  both  water  and  impurities.  When 
heated,  the  crystals  fly  to  pieces  explosively  (decrepitate),  on  account  of 


POTASSIUM   AND   AMMONIUM 


557 


the  vaporization  of  this  water.  All  substances  which  form  large  crystals 
and  do  not  melt  when  warmed,  behave  in  the  same  way  and  for  the  same 
reason.  In  consequence  of  this,  the  purest  salt  is  made  by  violent  stir- 
ring of  the  solution  during  the  operation  of  crystallization,  the  result 
being  the  formation  of  a  crystal-meal. 

Potassium  nitrate  is  used  chiefly  in  the  manufacture  of  gunpowder, 
which  contains  75  per  cent  of  the  highly  purified  salt.  The  other  com- 
ponents are  10  per  cent  of  sulphur,  14  per  cent  of 
charcoal,  and  about  1  per  cent  of  water.  The  ingre- 
dients are  intimately  mixed  in  the  form  of  paste, 
and  the  material  when  dry  is  broken  up  and  sifted, 
grains  of  different  sizes  being  used  for  different  pur- 
poses. The  chemical  action  which  takes  place  when 
gunpowder  is  fired  in  an  open  space  probably  results 
chiefly  in  the  formation  of  potassium  sulphide,  car- 
bon dioxide,  and  nitrogen  : 

2KN03  +  30  +  S  ->  K2S  +  3C02  +  Na. 

The  explosion  occurring  in  firearms  follows  a  much 
more  complex  course,  and  half  of  the  solid  product 
is  said  to  be  potassium  carbonate.  The  pressure,  at 
the  temperature  of  the  explosion,  if  the  gases  could 
be  confined  within  the  volume  originally  occupied  by  the  gunpowder, 
would  reach  about  forty-four  tons  per  square  inch.  In  recent  years 
common  gunpowder  has  been  displaced  largely  by  smokeless  powder, 
of  which  substances  related  to  gun-cotton  (p.  441)  are  the  chief  compo- 
nents. 

Potassium  Carbonate  K2CO3-  —  This  salt  is  manufactured 
from  potassium  chloride,  which  is  heated  with  magnesium  carbonate 
(magnesite),  water,  and  carbon  dioxide  under  pressure : 

2KC1  +  3MgC03  +  C02  +  5H20->2KHMg(C03)2,4H20  +  MgCL,. 

The  hydrated  mixed  salt  is  separated  from  the  liquid  containing 
magnesium  chloride  and  decomposed  by  heating  with  water  at 
120°.  The  product  is  a  solution  of  potassium  carbonate,  from  which 
the  precipitated  magnesium  carbonate  is  removed  by  filtration.  A 
certain  amount  is  also  obtained  from  the  fatty  material,  known  as 
suint,  which  forms  about  50  per  cent  of  the  weight  of  sheep's 
wool.  The  suint  is  separated  from  the  latter  by  washing.  When 


FIG.  98. 


558  INORGANIC    CHEMISTRY 

this  material,  which  contains  the  potassium  salt  of  sudoric  acid  in 
large  proportions,  is  calcined,  potassium  carbonate  remains,  and  is  ex- 
tracted from  the  ash  with  water.  Some  plants,  like  the  sugar-beet, 
take  up  exceptional  quantities  of  potassium  salts  from  the  soil.  The 
molasses  remaining  from  the  crystallization  of  beet-sugar  (p.  500) 
is  mixed  with  yeast  and  fermented.  After  the  alcohol  has  been  dis- 
tilled off,  the  liquid,  containing  organic  salts  of  potassium  in  solution, 
is  evaporated,  and  the  residue  is  ignited.  In  some  districts  potassium 
carbonate  is  still  extracted  from  wood-ashes. 

This  salt  is  usually  sold  in  the  form  of  an  anhydrous  powder  (m.-p. 
over  1000°).  When  crystallized  from  water  it  gives  a  hydrate  2K2C03, 
3H2O.  It  is  extremely  deliquescent.  Its  aqueous  solution  has  a 
marked  alkaline  reaction.  The  hydrolysis  of  the  salt  by  the  water  is 
exactly  analogous  to  that  of  sodium  sulphide  (p.  375),  although  not  so 
extensive.  The  more  elaborate  scheme  given  in  that  connection  may 
be  put  in  simpler  form  to  show  that  the  action  consists  essentially  in 
the  formation  of  the  ion  HCO8',  by  union  of  the  ion  C08"  with  the 
hydrion  of  the  water.  This  takes  place  because  the  ionization  of  the 
former  ion  is  small  enough  to  be  commensurable  with  that  of  water  it- 
self :  CO3"  +  H*  +  OH'  —  >  HCO/  -f  OH'.  The  commercial  name  of  the 
substance  is  pearl  ash.  It  is  used  in  making  soft  soap  and  hard  glass. 
It  is  also  employed,  by  interaction  with  acids,  in  making  salts  of 
potassium. 

When  a  concentrated  solution  of  the  salt  is  electrolyzed  in  such  a 
way  that  the  anode,  towards  which  the  KC03'  ions  travel,  consists  of 
a  thin  platinum  wire,  the'  crowding  together  of  the  discharged  material 
results  in  the  formation  of  the  percarbonate  (cf.  p.  397)  : 


The  operation  must  be  conducted  between  —  15°  and  0°.  When  the 
solution  in  the  porous  cell  surrounding  the  anode  is  evaporated,  the 
product  is  obtained  as  an  amorphous  bluish-  white  powder.  The  sub- 
stance liberates  oxygen  when  heated,  and  in  other  respects  behaves  like 
the  persulphates.  When  it  is  treated  with  a  dilute  acid,  a  solution 
containing  hydrogen  peroxide  is  formed.  The  compound  is  therefore 
a  mixed  anhydride  (p.  397)  of  hydrogen  peroxide  and  potassium  bi- 
carbonate. 

Potassium  Cyanide.  —  Formerly  this  compound  was  made  by 
heating  potassium  carbonate  with   nitrogenous   animal   matter.      So 


POTASSIUM   AND   AMMONIUM  559 

many  other  substances  were  formed  at  the  same  time,  however,  that 
the  required  product,  which  is  very  soluble,  was  difficult  to  isolate  in 
a  state  of  purity.  It  is  now  made  by  heating  together  potassium  ferro- 
cyanide  (q.v.)  and  potassium  carbonate.  The  ferrocyanide  acts  as  if  it 
were  a  mixture  of  potassium  cyanide  and  ferrous  cyanide :  K4Fe(CN)6 
— ->  4KCN  -h  Fe(CN)2.  The  latter,  by  interaction  with  the  potassium 
carbonate,  would  give  potassium  cyanide  and  ferrous  carbonate,  but  this 
in  turn,  is  decomposed  by  heat  into  ferrous  oxide,  which  is  insoluble, 
and  carbon  dioxide : 

K4Fe(CN)6  +  K2C03  _»  6KCN  +  FeO  +  CO2. 

When  the  residue  is  extracted  with  water,  only  the  potassium  cyanide 
dissolves,  and  it  is  easily  crystallized  in  pure  form  from  the  solution. 
Very  interesting  is  the  formation  of  potassium  cyanide  in  the  blast 
furnace  (q.v.).  Carbon  and  nitrogen  unite  at  a  very  high  temperature 
to  form  cyanogen  (p.  260),  and  a  sufficient  amount  of  potassium  is 
found  in  the  materials  to  complete  the  production  of  the  salt. 

Potassium  cyanide  crystallizes  in  cubes.  It  is  extremely  soluble  in 
water,  and  is  therefore  deliquescent.  Its  poisonous  qualities  are  equal 
to  those  of  hydrocyanic  acid.  The  acid  is  so  feeble  as  to  be  liberated 
even  by  the  carbon  dioxide  of  the  air,  and  hence  this  salt  always^  has  a 
distinct  odor  of  hydrocyanic  acid.  Potassium  cyanide  has  a  great 
tendency  to  form  complex  compounds  with  cyanides  of  other  metals 
(cf.  p.  536).  Complex  compounds  of  this  kind  are  used  in  the  gal- 
vanic deposition  of  silver  and  gold  in  commercial  electroplating. 
Large  amounts  of  the  cyanide' are  also  used  in  extracting  gold  (q.v.)  from 
its  ores,  particularly  in  the  Transvaal  colony.  The  tendency  to  form 
complex  compounds  is  doubtless  connected  with  the  fact  that  the 
cyanides  are  unsaturated  compounds  in  which  the  carbon  has  two 
free  valences  :  K— N  =  C  (p.  507). 

Potassium  cyanate  KCETO  is  made  by  heating  potassium  cyanide 
in  the  air,  or,  still  better,  with  some  easily  decomposed  oxide  (p.  507). 
It  is  a  white,  easily  soluble,  crystalline  salt. 

Potassium  thiocyanate  KCNS  may  be  obtained  by  melting  potas- 
sium cyanide  with  sulphur  (cf.  p.  508).  It  is  a  white,  deliquescent  salt 
which  finds  some  applications  in  chemical  analysis. 

The  Sulphate  and  Bisulphate.  —  The  sulphate  of  potassium 
is  a  constituent  of  several  double  salts  found  in  the  Stassfurt  de- 
posits. It  is  extracted  from  schoenite  MgS04,  KaS04, 6H?0  and  kainite 


560  INORGANIC  CHEMISTRY 

MgS04,  MgCl2,  K2SO4, 6H20.  The  former  is  treated  with  potassium 
chloride  and  comparatively  little  water,  whereupon  the  relatively  in- 
soluble potassium  sulphate  crystallizes  out,  and  the  magnesium  chlo- 
ride remains  in  the  mother-liquor.  The  crystals  belong  to  the  rhombic 
system,  contain  no  water  of  crystallization,  and  melt  at  1066°.  This  salt 
is  employed  in  large  quantities  in  making  potassium  carbonate  by  the  Le 
Blanc  process  and  in  preparing  alum  (q.v.).  It  is  also  much  used  as  a 
fertilizer.  Since  plants  take  up  solutions  through  their  cell  walls,  they 
can  absorb  soluble  compounds  only.  They  are,  therefore,  dependent, 
for  the  potassium  compounds  which  they  require,  upon  the  weathering 
out  of  soluble  potassium  compounds  from  the  insoluble  potassium 
silicates  contained  in  the  soil.  The  weathering  takes  place  too  slowly 
to  furnish  a  sufficient  supply  for  many  crops,  particularly  that  of  the 
sugar-beet.  Hence  potassium  sulphate  is  mixed  directly  with  the  soil. 
The  mineral  kainite  itself  is  used  for  the  same  purpose. 

Potassium  hydrogen  sulphate  (bisulphate)  KHS04  is  made  by 
the  action  of  sulphuric  acid  upon  potassium  sulphate :  K2S04  + 
H2SO4— >2KHSO4.  It  crystallizes  from  water,  in  which  it  is  very 
soluble,  in  tabular  crystals.  When  heated  to  about  200°  it  melts,  and 
the  elements  of  water  are  eliminated,  the  pyrosulphate  remaining  : 
2KH§O4  ->  H20  +  K2S2O7.  The  latter,  when  still  further  heated, 
yields  sulphur  trioxide  and  potassium  sulphate.  The  bisulphate  is 
used  in  analysis  for  the  purpose  of  decomposing  oxides  and  silicates 
and  converting  them  into  sulphates.  The  substance  is  more  efficient 
than  sulphuric  acid  for  this  purpose,  because  the  latter  cannot  be 
heated  above  330°,  while  the  liberation  of  the  active  sulphur  trioxide 
from  this  salt  takes  place  at  a  bright-red  heat.  The  aqueous  solu- 
tion of  the  bisulphate  is  strongly  acid  on  account  of  the  considerable 
ionization  of  the  hydrosulphanion. 

Sulphides  of  Potassium.  —  By  the  treatment  of  a  solution  of 
potassium  hydroxide  with  excess  of  hydrogen  sulphide,  a  solution  of 
potassium  hydrogen  sulphide  is  obtained.  Evaporation  of  the  solution 
gives  a  deliquescent  solid  hydrate  2KHS,  H20.  When  the  solution, 
before  evaporation,  is  treated  with  an  equivalent  amount  of  potassium 
hydroxide,  and  the  water  is  driven  off,  the  sulphide  K2S  remains 
behind  (cf.  p.  375) : 

KHS  +  KOH  <-»  K2S  +  H20. 

With  proper  care,  the  very  soluble  hydrate  K2S,  5H2O  may  be  obtained. 
Considerable  amounts  of  sulphur  can  be  dissolved  in  solutions  of  either 


POTASSIUM  AND  AMMONIUM  561 

of  thsse  sulphides.  By  evaporation  of  the  resulting  yellow  liquids,  vari- 
ous polysulphides  have  been  obtained.  To  some  of  these  have  been 
ascribed  the  formulae  K2S3,  K4S7,  K2S4,  K4S9,  and  K2S5  (cf.  p.  376). 
Similar  substances  are  produced,  as  a  result  of  the  liberation  and 
recombination  of  sulphur,  when  the  solutions  are  exposed  to  the 
oxidizing  action  of  the  air  : 

2KHS  +  0  ->  2KOH  +  2S. 


In  most  respects  the  corresponding  compounds  of  potassium  and 
sodium  are  similar  in  their  physical  properties  and  chemical  action. 
Since,  however,  the  latter  are  almost  uniformly  less  expensive,  they 
find  much  wider  application.  In  a  few  cases,  however,  the  potassium 
salt  is  more  generally  used.  Thus,  potassium  chlorate  and  potassium 
iodide  are  much  less  soluble  than  the  corresponding  sodium  compounds, 
and  it  is  consequently  possible  in  each  of  these  two  cases  to  separate 
by  crystallization,  and  to  purify  the  potassium  salt  with  greater  ease. 

Properties  of  Kalion  :  Analytical  Reactions.  —  The  positive 
ionic  material  of  the  potassium  salts  is  a  colorless  substance.  It 
unites  with  all  negative  ions,  and  most  of  the  resulting  compounds  are 
fairly  soluble.  For  its  recognition  we  add  solutions  containing  those 
ions  which  give  with  it  the  least  soluble  salts.  Thus,  with  chloroplat- 
inic  acid  H2PtCl6  it  gives  a  yellow  precipitate  of  potassium  chloro- 
platinate  K2PtCl6.  Since  nearly  one  part  of  this  salt  dissolves  in  100 
parts  of  water,  the  test  is  far  from  being  a  delicate  one.  The  solubility 
in  alcohol  is  much  smaller,  and  consequently  the  precipitate  may  fre- 
quently be  obtained  from  a  dilute  solution  by  adding  more  than  an 
equal  volume  of  alcohol.  Picric  acid  (p.  441)  gives  potassium  picrate 
KC6H2(1S~02)30,  which  is  much  less  soluble  in  water  (0.4  parts  in  100 
at  15°).  Perchloric  acid  and  hydrofluosilicic  acid  likewise  give  some- 
what insoluble  salts  of  potassium.  Potassium  hydrogen  tartrate 
KHC4H4O6  is  precipitated  by  the  addition  of  tartaric  acid  to  a  suffi- 
ciently concentrated  solution  of  a  potassium  salt.  The  normal  tartrate 
K2C4H4O6  is  much  more  soluble.  It  may  be  obtained  by  treating  the 
precipitate  with  a  solution  of  potassium  carbonate  or  potassium  hy- 
droxide. Addition  of  an  acid  to  this  solution  causes  reprecipitation 
of  the  bitartrate. 

The  Spectroscope.  —  A  much  more  delicate  test  for  the  recognition 
of  a  potassium  compound  consists  in  the  examination  by  means  of  the 


562  INORGANIC   CHEMISTRY 

spectroscope  of  the  light  given  out  by  a  Bunsen  flame,  in  which  a 
little  of  the  salt  is  held  upon  a  platinum  wire.  When  the  amount  of 
potassium  is  considerable,  and  no  other  substance  which  would  likewise 
color  the  flame  is  present  to  mask  the  effect,  the  violet  tint  is  recog- 
nizable by  the  eye.  In  general,  however,  the  light  must  be  analyzed. 

White  light  is  composed  of  vibrations  of  every  wave-length  within  a 
certain  range.  If  the  light  is  made  up  of  one  or  more  wave-lengths 
only,  it  appears  to  the  eye  to  be  colored.  Now,  when  a  narrow  bundle 
of  rays  of  white  light,  coming  through  a  slit,  falls  upon  a  three-sided 
prism  standing  with  its  edges  parallel  to  the  slit,  the  rays  of  various 
wave-length  are  retarded  to  different  extents  as  they  pass  through  the 
glass,  and  in  consequence  are  bent  from  their  paths  by  varying  amounts. 
Fig.  99  shows  a  horizontal  section  through  the  slit  (S)  and  prism,  in 

which  the  width  of  the  slit  and 
of  the  beam  of  light  are  exagger- 
Vioiet  ated.  The  light  emerging  at  the 
Blue  other  side  of  the  prism  consists, 
Yellow  therefore,  of  a  series  of  images 
Red  of  the  slit  arranged  side  by  side. 
The  red  light  is  least  refracted, 
and  the  red  images  of  the  slit, 
therefore,  are  most  nearly  in  the 
same  straight  line  with  the  ori- 
ginal beam.  The  yellow,  green, 

blue,  and  violet  images  are  displaced  more  and  more  from  this  direction, 
and  the  resulting  colored  band  is  called  a  spectrum.  The  whole 
series  of  images  of  the  slit  may  be  received  upon  a  screen,  or  directly 
upon  an  eye  looking  towards  the  prism.  Now,  when  the  light  comes 
from  the  vapor  of  potassium  heated  in  a  Bunsen  flame,  there  are  pro- 
duced, not  thousands  of  images  of  the  slit,  representing  as  many  differ- 
ent wave-lengths  of  light,  but  only  two  images,  one  red,  and  one  deep 
blue,  corresponding  to  the  two  wave-lengths  which  are  alone  contained 
in  the  original  light.  In  a  more  powerful  instrument  other  fainter 
lines  are  seen  also.  Naturally  the  brightness  of  all  these  lines  is 
together  equal  to  that  of  the  original  beam.  No  other  substance  gives 
any  of  those  particular  lines,  although  many  others  give  blue  and  red 
light  of  somewhat  different  wave-lengths.  Thus,  strontium  compounds 
give  a  blue  light  along  with  several  red  tints,  but  when  strontium  and 
potassium  are  used  together,  the  lines  are  found  not  to  be  coincident. 
In  the  case  of  strontium,  all  the  lines  lie  nearer  to  the  yellow  than  in 


POTASSIUM  AND  AMMONIUM  563 

that  of  potassium.  Since  the  whole  light  of  the  compound  is  thus  con- 
centrated  in  one  or  two  narrow  strips  easily  visible  against  a  dark 
background,  small  amounts  of  the  elements  give  effects  which  are 
readily  recognizable  in  the  instrument.  This  remains  true  even  when, 
to  the  eye,  the  colors  are  completely  obscured  by  the  much  more  bril- 
liant, yellow  light  which  compounds  of  sodium  produce.  In  the  spec- 
trum of  sodium,  this  yellow  light  is  all  concentrated  into  two  yellow 
lines  which  lie  very  close  together. 

Helium  gives  many  lines,  but  one  orange  line  (D3),  in  particular,  was 
noted  in  the  spectrum  of  the  sun's  photosphere  many  years  before  the 
element  was  obtained  from  terrestrial  sources  by  Eamsay.  When  the 
spectra  of  helium  and  other  gases  are  to  be  examined  in  the  laboratory, 
a  little  of  the  material  is  inclosed  in  a  narrow,  exhausted  tube,  through 
which  an  electrical  discharge  can  be  passed  between  platinum  wires. 
Under  this  treatment  helium  shows  its  conspicuous  orange  line,  and 
hydrogen  a  red  and  two  blue  ones.  In  this  apparatus  compounds  are 
dissociated  and  give  the  spectra  of  their  constituents.  When  a  Bunsen 
flame  is  used  with  the  salts*  of  metals,  however,  the  temperature  is  not 
high  enough  to  render  visible  the  spectra  of  the  non-metals  contained 
in  them.  Indeed,  even  of  the  metals  themselves,  only  the  members 
of  the  alkali  and  alkaline-earth  groups  give  distinct  results. 

RUBIDIUM  AND  CAESIUM. 

Soon  after  the  invention  of  the  spectroscope  by  Bunsen  and  Kirch- 
hoff,  the  instrument  was  applied  to  the  examination  of  many  sub- 
stances. In  1860  Bunsen  discovered  several  new  lines  in  the  spectrum 
given  by  materials  derived  from  the  salts  in  Diirkheim  mineral  water. 
Two  new  elements  of  the  alkali  group  were  found  to  cause  their 
presence,  and  were  named,  from  the  colors  of  the  lines  which  they  gave, 
rubidium  (red)  and  caesium  (blue).  Both  elements  have  since  been 
found  in  small  quantities  in  various  minerals.  Kubidium  is  obtain- 
able with  relative  ease  from  the  mother-liquors  of  the  Stassfurt  works. 

The  metals  may  be  obtained  by  heating  their  hydroxides  with 
magnesium  powder.  The  salts  of  these  two  elements  are,  in  crystal- 
line form  and  solubility,  very  much  like  those  of  potassium.  In  some 
cases  the  difference  in  solubility  is  sufficient  to  make  separation  pos- 
sible. Thus,  a  mixture  containing  compounds  of  these  two  metals 
and  of  potassium  gives  with  chloroplatinic  acid  a  yellow  precipitate, 

*  The  chlorides  are  preferred  because  of  their  volatility.  The  salts  of  the  oxy- 
gen acids  are  dissociated,  and  leave  the  highly  involatile  oxides  (e.g.  pp.  380,  444) 


564  INORGANIC   CHEMISTRY 

consisting  of  the  three  insoluble  chloroplatinates.  The  solubilities  at 
10°,  however,  are  as  follows :  Potassium  chloroplatinate  0.9,  rubidium 
chloroplatinate  0.15,  caesium  chloroplatinate  0.05.  Hence,  when  the 
mixed  precipitates  are  carefully  washed  with  small  quantities  of  cold 
water  the  potassium  chloroplatinate  can  be  almost  entirely  removed. 
On  similar  principles  the  two  other  metals  can  be  separated  from  one 
another.  The  iodides  of  all  three  elements  combine  with  iodine,  giving 
tri-iodides  (cf.  p.  235),  of  which  the  tri-iodide  of  caesium  is  the  most 
stable.  Whether  this  is  to  be  regarded  as  showing  that  the  metals  may 
occasionally  be  trivalent,  or  whether  the  extra  iodine  must  be  held  to 
have  entered  into  combination  with  the  iodine  of  the  compound,  and  not 
with  the  metal,  has  not  been  determined.  In  the  parallel  case  of 
hydriodic  acid,  the  union  with  extra  iodine  (p.  359)  seems  to  show  con- 
clusively that  iodidion  can  combine  with  iodine.  While  an  inclination 
to  tri valence  in  one  of  the  metals  of  the  alkalies  would  furnish  a  very 
acceptable  link  between  the  two  sides  of  the  first  column  in  the  peri- 
odic table  (p.  411),  since  gold  is  a  trivalent  element,  the  latter  of  the 
two  above  assumptions  is  more  probably  the  correct  one. 

AMMONIUM. 

The  compounds  of  ammonium  claim  a  place  with  those  of  the 
alkali  metals  because  in  aqueous  solution  they  give  the  ion  NH4',  an 
ion  which  in  its  behavior  closely  resembles  kalion.  Some  of  the 
special  properties  peculiar  to  ammonium  compounds  have  been  dis- 
cussed in  detail  already  (pp.  420,  421). 

Ammonium  Chloride.  —  This  salt,  known  commercially  as 
salammoniac,  like  all  the  other  compounds  of  ammonium,  is  prepared 
from  the  ammonia  dissolved  by  the  water  used  to  wash  illuminating- 
gas  (p.  418).  It  is  purified  by  sublimation,  and  then  forms  a  com- 
pact fibrous  mass.  It  crystallizes  from  solution  in  cubes  or  octahe- 
drons, which  are  often  arranged  according  to  a  feathery  pattern. 
When  heated  to  350°  it  volatilizes  and  is  almost  completely  dissoci- 
ated into  ammonia  and  hydrogen  chloride  at  this  temperature  (p.  421). 

Ammonium  bromide  and  ammonium  iodide  are  white  salts  which 
crystallize  in  cubes  or  octahedrons,  and  are  isomorphous  with  the  cor- 
responding potassium  salts.  They  are  dissociated  by  heat,  and,  in 
the  case  of  the  iodide,  some  of  the  hydrogen  iodide  is  still  further 
decomposed,  giving  free  iodine. 


POTASSIUM   AND  AMMONIUM  565 

Ammonium  Hydroxide.  —  The  nature  and  behavior  of  this  sub- 
stance have  been  fully  discussed  (p.  421).  It  may  be  remarked  here 
that  its  very  small  basic  activity  as  compared  with  that  of  potassium 
hydroxide  is  only  in  part  due  to  the  low  degree  of  ionization  of  its 
molecules.  A  normal  solution  of  ammonia  contains  much  free  NH8, 
besides  the  NH4OH  produced"  by  its  union  with  water.  Thus,  some 
part  of  the  material  —  how  much  we  have  at  present  no  means  of  de- 
termining —  is  not  actually  in  the  form  of  a  base  and  is  not  in  directly 
ionizable  condition  at  all.  There  are  indications  that  the  amount  of 
uncombined  ammonia  may  be  considerable.  Thus  the  organic  deriva- 
tive tetramethylammonium  hydroxide  N(CH8)4OH  is  a  very  active 
base  indeed,  and  one  of  the  most  conspicuous  differences  between  it 
and  ammonium  hydroxide  is  that  it  cannot  decompose  into  water  and 
a  non-ionizable  substance.  It  is  all  available  for  ionization,  while  the 
material  in  ammonia-water  is  not. 

Ammonium  Nitrate.  —  This  is  a  white  crystalline  salt  which 
may  be  made  by  the  interaction  of  ammonium  hydroxide  and  nitric 
acid.  When  heated  gently  it  decomposes,  giving  nitrous  oxide  and 
water  (p.  450).  It  is  used  as  an  ingredient  in  fireworks  and  explo- 
sives. It  exists  in  no  fewer  than  four  solid  physical  states.  The 
melted  salt  solidifies  at  about  160°,  giving  crystals  of  the  regular 
system.  When  these  are  allowed  to  cool  somewhat,  and  are  held  at 
a  temperature  a  little  below  125.5°,  they  change  gradually  into  a  mass 
of  rhombohedral  crystals,  the  specific  gravity  and  all  other  physical 
properties  altering  at  the  same  time.  This  temperature  is  a  transition 
point  like  that  at  which  monoclinic  sulphur  assumes  the  rhombic  form 
(p.  368).  When  these  rhombohedral  crystals,  in  turn,  are  held  at  a 
temperature  a  little  below  83°  they  change  their  form  once  more  into 
crystals  which  belong  to  the  rhombic  system  and  possess  a  third  dis- 
tinct set  of  physical  properties.  Finally,  below  35°  a  fourth  change, 
into  rhombic  needles,  takes  place,  and  this  condition  of  the  substance 
is  the  one  familiar  at  ordinary  temperatures.  All  these  changes  pro- 
ceed in  the  reverse  order  when  the  temperature  is  elevated  once  more. 

Ammonium  Carbonate.  —  When  ammonium  hydroxide  is  treated 
with  excess  of  carbon  dioxide  the  solution  gives,  on  evaporation, 
ammonium  bicarbonate  NH4HC08.  This  is  a  white  crystalline  salt 
which  is  fairly  stable  at  the  ordinary  temperature.  It  has,  however, 
a  faint  odor  of  ammonia,  and  its  dissociation  becomes  very  rapid  when 


566  INORGANIC  CHEMISTRY 

slight  heat  is  applied.     When  a  solution  of  this  salt  is  treated  with 
ammonium  hydroxide,  the  neutral  carbonate  is  formed : 

NH4HC03  +  NH4OH  <=±  (NH4)2C03  +  H20. 

But  this  salt,  when  left  in  an  open  vessel,  loses  ammonia  very  rapidly, 
and  leaves  the  bicarbonate  behind. 

The  substance  commonly  sold  as  ammonium  carbonate  is  the  so- 
called  aesquicarbonate,  and  is  made  by  sublimation  from  a  mixture  of 
ammonium  chloride  or  ammonium  sulphate  and  chalk  or  powdered 
limestone.  It  is  a  mixture,  in  approximately  equi-molar  proportions, 
of  ammonium  bicarbonate  and  ammonium  carbamate.  The  latter  is  a 
substance  related  to  urea,  and  formed  when  ammonia  and  carbon 
dioxide  gases  are  mixed : 

2NH3  +  C02  ->  0  =  0 


Ammonium  cyanate  is  interesting  on  account  of  its  rapid  trans- 
formation, when  warmed,  into  urea  (p.  488).  Ammonium  thiocyanate 
NH4NCS  is  a  white  salt  which  finds  some  application  in  analysis. 

Ammonium  Sulphate.  —  This  is  a  white  salt,  crystallizing  in 
rhombic  prisms,  which  is  used  chiefly  as  a  fertilizer.  By  electroylsis 
of  a  concentrated  solution  of  the  bisulphate,  ammonium  persulphate, 
which  is  less  soluble,  is  formed  and  crystallizes  out  (cf.  p.  397). 

Sulphides  of  Ammonium.  —  When  gaseous  hydrogen  sulphide 
and  ammonia  are  mixed  in  equi-molar  proportions  and  compressed  or 
strongly  cooled,  ammonium  hydrogen  sulphide  NH4HS  is  formed  as 
a  crystalline  deposit  on  the  vessel.  In  an  open  vessel,  at  the  ordinary 
temperature,  this  solid  dissociates  slowly  into  its  constituents.  The 
sulphide,  (NH4)2S,  can  be  produced  under  similar  conditions  by  using 
twice  as  much  ammonia.  But  it  is  much  less  stable  and  gives  up  half 
its  ammonia,  producing  the  acid  sulphide  very  quickly.  Solutions  of 
these  sulphides,  made  by  passing  hydrogen  sulphide  gas  into  ammo- 
nium hydroxide,  are  much  used  in  analysis.  The  sulphide  is  almost 
completely  hydrolyzed  by  water  into  the  acid  sulphide  and  ammonium 
hydroxide,  its  behavior  being  like  that  of  sodium  sulphide  (p.  375)  : 

2NH3  +-  H2S  «=±  (NH4)2S  ±=>  2NH4*  +  S" 

H0«--     OH'+F-^HS/ 

X10V_J  ^3         \Jtt.     -+•    -L7 


POTASSIUM  AND   AMMONIUM  567 

It  is  used  for  the  precipitation  of  sulphides,  such  as  zinc  sulphide, 
which  are  insoluble  in  water.  Although  the  S"  ions  are  not  numerous 
at  any  moment,  disturbance  of  the  equilibrium  by  their  removal,  when 
they  pass  into  combination,  causes  displacements  which  result  in  the 
generation  of  a  continuous  supply.  The  liquid  smells  strongly  of 
ammonia  and  hydrogen  sulphide  on  account  of  the  dissociation  of  the 
parent  molecules.  Because  of  this  dissociation  the  salt  is  preferred  to 
potassium  or  sodium  sulphide  in  analysis.  The  excess  of  the  reagent 
can  be  driven  out  by  simply  boiling  the  mixture  for  a  few  minutes,  all 
of  the  above  equilibria  being  reversed.  Another  application  in  analy- 
sis depends  on  the  tendency  of  this  salt  to  unite  with  certain  insoluble 
sulphides,  particularly  those  of  tin,  arsenic,  and  antimony  (q.v.\  giving 
soluble  complex  salts. 

The  solution  dissolves  free  sulphur,  giving  yellow  polysulphides 
similar  to  those  of  potassium  (p.  561).  The  same  yellow  substances 
are  also  obtained  by  gradual  oxidation  of  ammonium  sulphide  when  the 
solution  of  this  salt  is  allowed  to  stand  in  a  bottle  from  which  the  air  is 
imperfectly  excluded. 

Microcosmic  Salt.  —  This  salt  would  be  named,  systematically, 
the  tetrahydrate  of  secondary  sodium-ammonium  orthophosphate 
(NaNH4HP04,  4H20).  When  ammonium  chloride  and  ordinary  sodium 
phosphate  are  mixed  in  strong  solution  the  hydrate  crystallizes  out. 
The  substance  is  used  in  bead  tests  (cf.  pp.  467,  468). 

Ammonium  Amalgam.  —  As  we  have  seen  (p.  550),  a  potassium 
ion  may  be  discharged  and  the  element  which  it  contains  secured  in  the 
free  condition.  When  a  salt  of  ammonium  is  decomposed  by  electro- 
lysis, however,  the  NH4  ion  upon  its  discharge  gives  ammonia  and  hy- 
drogen, and  no  substance  NH4  is  obtained.  If,  however,  a  pool  of 
mercury  is  used  as  the  negative  electrode,  the  NH4  forms  an  amalgam 
with  it,  and  there  seems  to  be  no  doubt  that  this  substance  is  actually 
present  in  solution  in  the  mercury.  While  the  amalgam  is  being  formed 
it  swells  up  and  gives  off  the  decomposition  products  above  mentioned, 
so  that  the  existence  of  the  substance  is  only  temporary.  The  same 
material  may  be  obtained  by  putting  sodium  amalgam  into  a  strong 
solution  of  a  salt  of  ammonium.  The  action  is  a  displacement  of 
one  ion  by  another  (p.  361) : 

Na  (diss'd  in  mercury)  -f  NH4*  — »  NH4  (diss'd  in  mercury)  -f  Na". 


568  INOEGANIC  CHEMISTRY 

This  behavior  is  interesting  since  it  is  in  harmony  with  the  idea  that 
ammonium,  if  it  could  be  isolated,  would  have  the  properties  of  a  metal. 
Substances,  other  than  metals,  are  not  miscible  with  mercury. 

Ammonion :  Analytical  Reactions. — Ionic  ammonium  is  a  color- 
less substaMce.  It  unites  with  negative  ions,  giving  salts,  which,  in  the 
majority  of  cases,  are  soluble.  Ammonium  chloroplatinate,  and  to  a 
less  extent  ammonium  hydrogen  tartrate,  are  insoluble  compounds,  and 
their  precipitation  is  used  as  a  test.  The  surest  means  of  recognizing 
ammonium  compounds,  however,  consists  in  adding  a  soluble  base  to 
the  substance  (cf.  p.  421).  The  ammonium  hydroxide,  which  is  thus 
formed,  gives  off  ammonia,  and  the  latter  may  be  detected  by  its  odor. 
The  quantity  of  the  ammonium  salt  present  may  be  determined  by  dis- 
tilling the  mixture  and  catching  the  distillate  in  a  measured  volume  of 
normal  hydrochloric  acid.  Determination  of  the  amount  of  the  acid 
remaining  unneutralized,  by  titration  with  a  standard  alkali  solu- 
tion, then  gives,  by  difference,  the  quantity  of  ammonium  hydroxide. 

Exercises.  —  1.  What  kind  of  metals  will,  in  general,  interact  with 
solutions  of  bases  (cf.  p.  553)  ? 

2.  Why  should  a  mixture  of  potassium  chlorate  and  antimony  tri- 
sulphide  be  explosive  ? 

3.  How  does  the  direct  vision  spectroscope  differ  from  the  arrange- 
ment here  described  (cf.  any  work  on  physics)  ? 

4.  Why  is  not  ammonium  carbamate  (p.  566)  formed  by  the  neutral- 
ization method? 

5.  How  should  you  set    about   making  :  a  borate  of   potassium, 
potassium  pyrophosphate,  ammonium  nitrite  ? 

6.  Why  is  the  cleaning  of  platinum  wires,  as  usually  effected  by 
holding  them   in   the   Bunsen  flame,  assisted  by  periodical   dipping 
into  hydrochloric  acid  (p.        )  ? 


CHAPTER  XXXIV 

SODIUM    AND    LITHIUM.      IONIC  EQUILIBRIUM   CONSIDERED 
QUANTITATIVELY 

SODIUM  chloride  forms  more  than  two-thirds  of  the  solid  matter 
dissolved  in  sea-water,  and  the  great  salt  deposits  are  largely  composed 
of  it.  Sea-plants  contain  sodium  salts  of  organic  acids,  just  as  land- 
plants  contain  potassium  salts.  Chili  saltpeter,  cryolite,  and  albite  (a 
soda  feldspar)  are  important  minerals. 

Preparation.  —  Sodium  was  first  made  by  Davy  (1807)  by  electro- 
lysis of  moist  sodium  hydroxide.  It  is  manufactured  by  Castner's 
process,  which  is  used  also  for  potassium  (p.  550),  and  by  the  electro- 
lysis of  fused  sodium  hydroxide  by  a  method  likewise  invented  by 
Castner.  In  the  latter  case  the  negative  electrode  projects  through  the 
bottom  of  the  iron  vessel  containing  the  fused  hydroxide.  This  electrode 
is  surrounded  by  a  wire-gauze  partition,  which  is  surmounted  by  a  bell- 
shaped  vessel  of  iron.  The  positive  electrode  is  an  iron  cylinder  sur- 
rounding the  gauze.  The  sodium  and  hydrogen  liberated  at  the  cathode, 
being  lighter  than  the  fused  mass,  ascend  into  the  iron  vessel,  under 
the  edge  of  which  the  hydrogen  escapes.  Oxygen  is  set  free  at  the 
anode. 

Properties.  —  Sodium  is  a  soft,  shining  metal,  melting  at  95.6° 
and  boiling  at  742°.  The  vapor  is  a  monatomic  gas.  The  metal  is 
soluble  in  liquefied  ammonia,  giving  a  blue  solution.  The  amalgam  with 
mercury,  when  it  contains  more  than  a  small  amount  of  sodium,  is 
solid,  and  probably  contains  one  or  more  compounds  of  the  two  elements. 
This  amalgam  is  often  used  instead  of  the  metal  sodium,  since  the 
dilution  or  combination  with  mercury  makes  the  interactions  of  the 
metal  more  easily  controllable.  Sodium  is  used  in  the  manufacture  of 
many  complex  carbon  compounds  which  are  employed  as  drugs  and  dyes. 

Sodium  Hydride.  —  When  hydrogen  is  led  over  sodium  at  340° 
in  such  a  way  that  the  upper  part  of  the  tube  is  cooler,  a  matted  mass 
of  fine  white  crystals  of  the  hydride  is  deposited  on  the  cool  part  of 

669 


570  INORGANIC   CHEMISTRY 

the  tube.  The  temperature  must  not  rise  beyond  430°,  since  the  com- 
pound  dissociates  rapidly  at  this  temperature.  The  properties  of  the 
substance  are  similar  to  those  of  potassium  hydride  (p.  550). 

Sodium  Chloride.  —  Common  salt  is  obtained  from  the  salt  de- 
posits of  Stassfurt,  Reichenhall  (near  Salzburg),  in  Cheshire,  at  Syra- 
cuse and  Warsaw  in  New  York,  at  Salina  in  Kansas,  in  Utah,  Cali- 
fornia, and  many  other  districts.  Natural  brines  are  obtained  from 
wells  in  various  parts  of  the  world.  Since  the  salt  can  seldom  be 
used  directly,  on  account  of  impurities  which  it  contains,  it  is  purified 
by  recrystallization  from  water.  Natural  brines,  which  are  sometimes 
dilute,  are  often  concentrated  by  dripping  over  extensive  ricks  com- 
posed of  twigs.  When  the  resulting  brine  is  allowed  to  evaporate 
slowly  by  the  help  of  the  sun's  heat,  large  crystals,  sold  as  ."solar 
salt,"  are  obtained.  By  the  use  of  artificial  heat  and  stirring,  smaller 
crystals  of  greater  purity  can  be  secured.  Salt  intended  for  table 
use  must  be  freed  from  the  traces  of  magnesium  chloride  (y.v.)  present 
in  the  original  brine  or  deposit,  for  this  impurity  causes  it  to  absorb 
moisture  more  vigorously  from  the  air.  The  purest  salt  for  chemical 
purposes  is  precipitated  from  a  saturated  solution  of  salt  by  leading 
into  it  hydrogen  chloride  gas.  Explanation  of  this  effect  will  be 
given  presently  (see  p.  584). 

Common  salt  crystallizes  in  cubes,  the  faces  of  which  are  usually 
hollow.  The  crystals  decrepitate  (p.  556)  when  heated,  and  melt  at 
about  820°.  Common  salt  is  the  source  of  all  sodium  compounds,  with 
the  exception  of  the  nitrate.  From  it  come  also  most  of  the  chlorine 
and  hydrogen  chloride  used  in  commerce. 

The  Hydroxide  and  Oxides.  —  Sodium  hydroxide  is  prepared 
both  by  the  action  of  slaked  lime  upon  sodium  carbonate  and  by  the 
electrolysis  of  a  solution  of  sodium  chloride,  precisely  as  is  potassium 
hydroxide  (p.  554).  An  interesting  method  used  at  Niagara  Falls, 
the  Acker  process,  employs  molten  sodium  chloride  as  the  electrolyte 
and  molten  lead  as  the  cathode.  The  sodium  forms  an  alloy  with  the 
lead,  which  continually  circulates  along  the  trough-shaped  cell  and 
back  beneath  a  partition.  On  issuing  at  one  end  it  comes  in  contact 
with  a  jet  of  steam,  and  this  removes  the  sodium  from  the  lead,  forming 
sodium  hydroxide.  The  latter  flows  in  the  molten  condition  from  the 
surface  of  the  lead,  and  the  metal  returns  to  the  trough. 

Sodium  hydroxide  is  a  highly  deliquescent  substance,  which,  when 


SODIUM  571 

exposed  to  the  air,  first  liquefies  and  then  becomes  solid  on  account  of 
the  formation  of  sodium  carbonate.  Its  general  chemical  properties 
are  identical  with  those  of  potassium  hydroxide.  It  is  used  in  the 
manufacture  of  soap,  in  the  preparation  of  paper  pulp,  and  in  many 
other  chemical  industries. 

Sodium  peroxide  Na202  is  made  by  passing  slices  of  sodium,  rest- 
ing upon  trays  of  aluminium,  through  a  tubular  vessel.  In  this  it  comes 
in  contact  with  a  current  of  air  which  has  been  freed  from  carbon 
dioxide  and  is  maintained  at  a  temperature  between  300°  and  400°. 
This  oxide  when  thrown  into  water  decomposes  in  part,  in  consequence 
of  the  heat  developed,  giving  sodium  hydroxide  and  oxygen.  With 
careful  cooling,  however,  much  of  it  can  be  dissolved.  By  interaction 
with  acids  it  yields  hydrogen  peroxide.  Sodium  peroxide  is  now  used 
commercially  for  oxidizing  and  bleaching.  The  ordinary  sodium  oxide 
Na20  is  made  in  the  same  way  as  is  potassium  oxide  (p.  555). 

The  Nitrate  and  Nitrite.  —  The  occurrence  and  purification  of 
sodium  nitrate  have  already  been  described  (p.  438).  Its  crys- 
tals are  of  rhombohedral  form  (Fig.  9,  p.  14).  This  salt  is  one 
of  the  best  of  fertilizers,  since  it  furnishes  to  plants  the  nitrogen 
which  they  require  in  a  very  easily  absorbed  form.  It  is  used  also  in 
the  manufacture  of  potassium  nitrate,  of  nitric  acid,  and  of  sodium 
nitrite. 

Sodium  nitrite  is  formed  by  heating  sodium  nitrate  with  metallic 
lead  and  recrystallizing  the  product  (p.  449).  Although  very  soluble 
it  is  less  so  than  potassium  nitrite,  and  is  therefore  more  easily  pre- 
pared in  pure  condition.  It  is  used  as  a  source  of  nitrous  acid  by 
manufacturers  of  organic  dyes. 

Manufacture  of  Sodium  Carbonate.  —  Natural  sodium  carbon- 
ate is  found  in  Egypt  and  in  other  parts  of  the  world.  At  Owen's 
Lake,  California,  it  is  secured  by  solar  evaporation  of  the  water. 
The  sesquicarbonate  Na2C08,NaHCO8,2H2O,  being  the  least  soluble 
of  the  carbonates  of  sodium,  is  the  one  deposited.  Locally,  small  quan- 
tities of  sodium  carbonate  are  still  made  by  the  burning  of  sea-weed. 
Up  to  the  close  of  the  eighteenth  century  this  was  the  only  source  of 
the  compound,  and  the  product  from  Spain,  known  commercially  as 
barilla,  was  ten  times  as  expensive  as  the  carbonate  now  is.  Hence 
glass  and  soap  were  proportionately  dearer  than  at  present. 

In  1791  the  French  Academy  offered  a  prize  for  the  discovery  of 


572 


INORGANIC   CHEMISTRY 


an  inexpensive  method  for  the  preparation  of  sodium  carbonate  from 
common  salt,  and  Le  Blanc  proposed  the  process  which  bears  his  name 
and  is  still  in  use.  During  the  Eevolution  his  factory  was  destroyed, 
his  patents  were  declared  to  be  public  property,  and  the  inventor  died 
by  suicide.  The  chief  stages  of  Le  Blanc's  process  involve  three  chemi- 
cal actions.  In  the  first  place,  sodium  chloride  is  treated  with  an  equi- 
valent amount  of  sulphuric  acid  in  a  large  cast-iron  or  earthenware 


FIG.  100. 

pan.  The  bisulphate  thus  produced  (cf.  p.  178),  together  with  the 
unchanged  sodium  chloride,  is  raked  out  on  to  the  hearth  of  a  rever- 
beratory  *  furnace  (Fig.  100)  and  heated  more  strongly,  while  being 
continually  worked  by  means  of  rakes,  until  the  action  is  completed  : 

NaCl  +  NaHS04  +±  Na2S04  +  HC1  1  .  ' 

The  product  of  this  treatment  is  called  salt-cake.  The  hydrogen 
chloride,  which  is  liberated  in  both  stages,  passes  through  towers 
containing  running  water  in  which  it  is  absorbed.  The  second  and 
third  actions  which  follow  are  conducted  in  one  operation.  They  con- 
sist in  the  reduction  of  the  sodium  sulphate  by  means  of  powdered 
coal  and  the  interaction  of  the  resulting  sulphide  of  sodium  with 
chalk  or  powdered  limestone  : 


+  20  - 
CaC03 


2CO 


Na2C03 


2, 

CaS. 


In  the  less  modern  factories  the  salt-cake,  limestone,  and  coal  are 
stirred  upon  the  hearth  of  a  reverberatory  furnace  and  worked  by  hand. 

*  So  called  because  the  heated  gases  from  the  fire  are  deflected  by  the  roof  and 
play  upon  the  materials  spread  on  the  bed  of  the  furnace. 


SODIUM  573 

The  material  is  finally  collected  into  balls,  and  the  end  of  the  action  is 
recognized  by  the  fact  that  bubbles  of  carbon  monoxide  begin  to  force 
their  way  to  the  surface  and  cause  little  jets  of  blue  flame.  The  gas 
is  produced  by  the  action  of  the  coal  upon  the  calcium  carbonate, 
excess  of  both  of  these  substances  being  present : 

CaC08  +  C  ->  CaO  +  2CO. 

The  production  of  this  gas  gives  a  porous  texture  to  the  material, 
which  facilitates  the  solution  of  the  sodium  carbonate  in  the  final 
stage.  The  porous  product  is  called  black-ash.  In  modern  factories 
hand  labor  is  saved  by  giving  the  black-ash  furnace  the  form  of  a 
rotating  cylinder,  in  which  projections  from  the  walls  assist  in  bring- 
ing about  complete  mixing  of  the  materials  during  the  action. 

The  black-ash  varies  very  much  in  composition.  It  commonly 
contains  45  per  cent  of  sodium  carbonate,  30  per  cent  of  calcium 
sulphide,  10  per  cent  of  calcium  oxide,  and  a  number  of  other  products 
and  impurities.  The  coal  used  in  the  operation  is  selected  so  as  to  be 
as  free  as  possible  from  combined  nitrogen,  the  presence  of  which 
leads  to  the  formation  of  cyanides. 

Calcium  sulphide  is  not  very  soluble  in  water,  and  is  but  slowly 
hydrolyzed  by  it  (p.  376),  especially  when  calcium  hydroxide  is  present. 
The  sodium  carbonate  is  therefore  extracted  from  the  black-ash  by  a 
systematic  treatment  of  the  ash  with  water.  The  ash  is  placed  in  a 
series  of  vessels  at  different  levels,  and  a  stream  of  water  flows  from 
one  vessel  to  another,  until,  when  it  issues  from  the  last,  it  is  com- 
pletely saturated  with  sodium  carbonate.  A  temperature  of  30°  to  40°, 
at  which  the  solubility  of  sodium  carbonate  is  at  a  maximum,  is 
employed.  When  the  material  in  the  first  of  the  vessels  has  been 
exhausted,  the  water  is  allowed  to  enter  the  second  vessel  directly,  and  a 
vessel  containing  fresh  black-ash  is  added  at  the  lower  end  of  the  series. 
In  this  way  the  most  nearly  exhausted  ash  comes  in  contact  with  pure 
water,  which  is  in  the  best  position  to  dissolve  the  remaining  sodium 
carbonate  rapidly,  while  the  fresh  black-ash  encounters  a  solution 
already  almost  at  the  point  of  saturation. 

The  saturated  solution  is  evaporated  in  shallow  pans  placed  in  the 
flues  of  the  furnaces,  and  the  monohydrate  Na^COg,  H2O,  which  crystal- 
lizes from  the  hot  liquid,  is  raked  out  and  dried  by  heat,  leaving  cal- 
cined soda.  When  this  material  is  recrystallized  from  water  and  is 
allowed  to  deposit  itself  from  the  solution  at  the  ordinary  temperature, 
the  decahydrate  Na2CO8,10H20,  soda  crystals  or  ivashing  soda,  appears. 


574  INORGANIC  CHEMISTRY 

The  solid  residue  from  the  extraction  of  the  black-ash  is  known  as 
tank  waste,  and  contains  35-55  per  cent  of  calcium  sulphide.  This 
material  contains  the  sulphur  of  the  original  sulphuric  acid,  and  its 
treatment  involves  a  problem  of  some  difficulty.  If  it  is  dumped  near 
the  factory  the  sulphur  is  lost,  and  by  slow  weathering  yellow  solutions 
containing  poly  sulphides  flow  from  the  decomposing  heap  into  the 
streams,  and  offensive  odors  of  hydrogen  sulphide  fill  the  air.  The 
most  effective  process  for  the  recovery  of  the  sulphur  and  consequent 
abatement  of  this  nuisance  is  that  of  Chance.  The  product  is  arranged 
in  a  series  of  cylinders  through  which  is  passed  carbon  dioxide  from 
a  kiln  of  special  form.  The  hydrogen  sulphide  liberated  in  the  first 
cylinder  forms  the  acid  sulphide  with  the  material  contained  in  the 
second : 

2CaS  +  C02  +  H20  ->  Ca(SH)2  +  CaC08. 

The  further  action  of  the  carbon  dioxide  on  this  product  gives,  finally, 
a  mixture  of  gases  containing  a  larger  proportion  of  hydrogen  sul- 
phide. By  burning  this  mixture  with  a  limited  supply  of  air,  the 
sulphur  is  then  secured  in  free  condition  : 

2H2S  +  02  ->  2H20  +  28. 

About  70,000  tons  of  sulphur  are  thus  recovered  annually. 

The  Solvay,  or  ammonia-soda  process,  invented  in  1860,  is  a  seri- 
ous rival  of  the  Le  Blanc  process.  It  differs  from  the  latter  by 
involving  almost  nothing  but  ionic  actions.  A  solution  of  salt  con- 
taining ammonia  and  warmed  to  40°  fills  a  tower  divided  by  a  number 
of  perforated  partitions.  Carbon  dioxide,  which  is  forced  in  below, 
makes  its  way  up  through  the  liquid.  The  ammonium  bicarbonate 
formed  by  its  action  undergoes  double  decomposition  with  the  salt,  and 
sodium  bicarbonate  which  is  precipitated  settles  upon  the  partitions  : 
NaCl+NH4HC03  <=>  NaHC03 J  +  NH4C1,  or  HC03'  +  Na'  <=±  NaHC08  J. 
The  solid  sodium  bicarbonate,  after  being  freed  from  the  liquid,  is 
heated  strongly  and  leaves  behind  sodium  carbonate  : 

2NaHC03  ->  NajCO,  +  H20  \  +  CO,  \. 

The  carbon  dioxide  which  is  liberated  passes  through  the  operation 
once  more.  The  mother-liquor  from  the  sodium  bicarbonate  contains 
ammonium  chloride.  This  is  decomposed  by  heating  with  quicklime, 
and  the  ammonia  which  is  thus  obtained  is  available  for  the  treatment 
of  another  batch.  The  supply  of  carbon  dioxide  is  generated  in 


SODIUM  575 

lime-kilns  of  special  form.  The  lime  produced  in  these  kilns  serves 
for  the  liberation  of  the  ammonia. 

This  process  is  cheaper  than  that  of  Le  Blanc,  and  furnishes  a 
much  purer  product.  The  latter  process  continues  to  be  used,  how- 
ever, because  of  the  hydrochloric  acid  which  is  produced  at  the  same 
time.  This  finds  remunerative  application  in  the  liberation  of  its 
chlorine  for  the  manufacture  of  bleaching  powder. 

A  possible  rival  of  these  two  processes  threatens  to  arise  in  the 
treatment  of  electrolytic  sodium  hydroxide  with  carbon  dioxide  gas. 

Properties  of  Sodium  Carbonate.  —  The  common  form  of 
sodium  carbonate  consists  of  large  monoclinic  crystals  of  the  decahy- 
drate.  This  substance  has  a  fairly  high  aqueous  tension,  and  loses 
nine  of  the  tea  molecules  of  water  which  it  contains  when  it  is  exposed 
in  an  open  vessel.  When  warmed  it  melts  at  35.2°,  giving  a  solution 
of  sodium  carbonate  in  water.  The  residue  from  evaporation,  above 
35.2°,  is  the  monohydrate.  At  higher  temperatures,  or  with  low  at- 
mospheric aqueous  tension  (p.  121),  this  in  turn  can  be  completely  dehy- 
drated (see  Chap.  xxxv).  The  decahydrate  increases  in  solubility  up 
to  35.2°,  when  it  ceases  to  exist.  Just  above  this  temperature  the  mono- 
hydrate  is  the  only  substance  which  is  stable.  Its  solubility  is  the 
same  as  that  of  the  decahydrate  at  35.2°,  and  diminishes  as  the  tem- 
perature is  raised.  The  relations  are  of  the  same  nature  as  in  the  case 
of  sodium  sulphate  (p.  158).  In  aqueous  solution,  sodium  carbonate  is 
hydrolyzed,  and  shows  a  marked  alkaline  reaction.  The  compound  is 
used  in  large  amounts  for  the  manufacture  of  glass  and  soap,  and 
is  applied  in  innumerable  ways  in  the  scientific  industries  for  pur- 
poses akin  to  cleansing. 

Nearly  all  the  familiar  compounds  of  sodium  are  formed  in  the 
course  of  one  or  other  of  the  processes  by  which  sodium  carbonate  is 
manufactured,  or  are  made  by  the  treatment  of  sodium  carbonate  or 
sodium  hydroxide  with  acids. 

Sodium  Bicarbonate.  —  This  salt  can  be  prepared  in  a  state  of 
purity  by  passing  carbon  dioxide  over  the  decahydrate  of  sodium 
carbonate : 

NasCO^lOHaO  +  Ca,^±2NaHC08  +  9H2O. 

The  hydrate  is  spread  upon  a  grating,  through  which  the  water  generated 
by  the  action  drips  away.  This  action  is  reversible,  and  sodium  bicar- 


576  INORGANIC   CHEMISTRY 

bonate  shows,  even  in  the  cold,  an  appreciable  tension  of  carbon  diox 
ide.  Even  a  solution  of  this  salt  gives  off  carbon  dioxide,  when  boiled. 
An  aqueous  solution  of  pure  sodium  bicarbonate  is  neutral  to  phenol- 
phthalein  on  account  of 'the  small  degree  of  ionization  of  the  ion  HC03'. 
The  salt  is  used  in  the  manufacture  of  baking  powder  and  in  medicine. 

Sodium  Sulphate.  —  Anhydrous  sodium  sulphate  (thenardite) 
crystallizes  in  the  rhombic  system,  and  is  found  in  the  salt  layers. 
The  same  salt  is  contained  in  mineral  waters,  such  as  those  of  Frie- 
drichshall  and  Karlsbad.  It  is  formed  in  connection  with  the  manu- 
facture of  nitric  acid  from  sodium  nitrate,  and  as  an  intermediate 
product  in  the  making  of  sodium  carbonate.  Some  of  it  is  also  pre- 
pared at  Stassfurt  by  dissolving  kieserite  (MgS04,H20)  in  water 
along  with  sodium  chloride.  When  the  solution  is  cooled  to  0°,  the 
decahydrate  of  sodium  sulphate  crystallizes  out,  and  magnesium  chlo- 
ride remains  in  solution. 

The  decahydrate  of  sodium  sulphate,  Glauber's  salt,  forms  large 
monoclinic  crystals  which  give  up  ten  molecules  of  water  when  kept  in 
an  open  vessel.  When  heated  the  crystals  melt  at  32.4°,  and  are  re- 
solved into  the  sulphate  and  water.  The  relations  of  the  hydrate  and 
anhydrous  substance  in  respect  to  solubility  have  been  fully  discussed 
already  (p.  158).  When  the  decahydrate  is  mixed  with  concentrated 
hydrochloric  acid,  it  is  decomposed,  and  a  part  of  the  sulphate  is  con- 
verted into  sodium  chloride,  the  second  action  being  a  reversible  one. 
This  is  one  of  those  actions  which  proceed  spontaneously,  and  there- 
fore involve  a  diminution  in  the  store  of  available  energy  in  the 
system,  although,  so  far  as  heat  is  concerned,  a  marked  absorption  of 
this  form  of  energy  takes  place  (cf.  p.  27).  The  combination  is  used, 
in  fact,  as  a  freezing  mixture. 

Sodium  Thiosulphate.  —  This  salt  is  made  by  boiling  a  solution  of 
sodium  sulphite  with  sulphur.  It  is  also  obtained  by  boiling  sulphur 
with  caustic  soda,  and  crystallizes  from  the  mixed  solution : 

4S  +  GNaOH  -»  Na,S,O,  +  2NaJ3  +  3H20. 

It  gives  large,  transparent  monoclinic  crystals  of  a  pentahydrate. 
When  heated  it  first  loses  the  water  of  hydration,  and  then  decomposes, 
giving  sodium  sulphate,  which  is  the  most  stable  oxygen-sulphur  com- 
pound of  sodium,  and  sodium  pentasulphide  : 

3  ->  3Na,S04 


LITHIUM  577 

From  the  latter,  four  unit- weights  of  sulphur  can  be  driven  by  stronger 
heating.  Sodium  thiosulphate  is  used  for  fixing  negatives  in  photog- 
raphy (q.v.),  and  by  bleachers  as  antichlor  (p.  396). 

Phosphates  of  Sodium.  —  Common  sodium  phosphate  is  a  do- 
decahydrate  of  the  secondary  orthophosphate,  Na2HP04,12H2O.  It  is 
made  by  neutralization  of  phosphoric  acid  with  sodium  carbonate,  and 
crystallizes  from  the  solution  in  large,  transparent  monoclinic  prisms. 
Its  properties  have  already  been  discussed  (pp.  466, 467). 

Sodium  metaphosphate  NaP08  is  used  for  bead  reactions  (cf.  p.  468). 

Sodium  Tetraborate.  —  This  salt  combines  with  ten  molecules  of 
water,  forming  large,  transparent  prisms,  Na2B4O7, 10H20.  When 
heated  it  loses  water,  and  leaves  the  easily  fusible  anhydrous  salt  in 
glassy  form.  Its  sources  have  already  been  discussed  under  boric  acid 
(p.  528).  It  is  used  as  an  ingredient  in  glazes  for  porcelain,  in 
soldering,  and  for  bead  reactions  (cf.  p.  528). 

Sodium  Silicate.  —  A  salt  essentially  of  the  composition  of  the 
metasilicate  NagSiOg  (cf.  p.  522)  is  used  for  fire-proofing  wood  and  other 
materials.  Sand  which  is  moistened  with  it  and  pressed  in  molds 
forms,  after  baking,  a  serviceable  artificial  stone.  Since  silicic  acid  is 
a  feeble  acid,  this  salt  is  much  hydrolyzed,  and  gives  a  strongly  alka- 
line solution  (p.  523). 

Properties  of  Natrion  :  Analytical  Reactions.  —  Natrion  is 
a  colorless  ionic  material  which  unites  with  all  negative  ions.  Practi- 
cally all  the  salts  so  formed  are  soluble  in  water.  The  only  ones  which 
can  be  precipitated  are  sodium  fluosilicate,  made  by  the  addition  of 
hydrofluosilicic  acid  to  a  strong  solution  of  a  sodium  salt,  and  sodium 
hydrogen  pyroantimoniate  Na2H2Sb207,  made  by  similar  addition  of  the 
corresponding  potassium  salt.  For  the  recognition  of  natrion  the 
solution  is  evaporated,  and  the  residue  examined  with  a  spectroscope. 
If  the  yellow  light  persists  longer  than  could  be  accounted  for  by  the 
ordinary  deposit  of  dust  on  the  wire,  a  sodium  compound  is  present  in 
the  material. 

LITHIUM. 

The  compounds  of  lithium  are  made  from  amblygonite,  a  mixed 
phosphate  and  fluoride  of  aluminium  and  lithium.  It  occurs  in  lepido- 
lite  (a  lithia  mica)  and  in  other  rare  minerals.  Traces  of  compounds 


578  INORGANIC  CHEMISTRY 

of  the  element  are  found  widely  diffused  in  the  soil,  and  are  taken  up 
by  plants,  particularly  tobacco  and  beets,  in  the  ashes  of  which  the 
element  may  be  detected  spectroscopically. 

The  metal  is  liberated  by  electrolysis  of  the  fused  chloride,  the 
manipulation  being  facilitated  by  the  addition  of  some  potassium  chlo- 
ride to  lower  the  melting-point  of  the  lithium  salt.  The  melting-point 
and  boiling-point  of  the  free  element  are  higher  than  those  of  any 
other  alkali-metal,  and  the  specific  gravity  (0.53)  is  lower  than  that 
of  any  other  metal  whatever.  Lithium  not  only  floats  upon  water, 
but  also  in  the  petroleum  in  which  it  is  preserved. 

The  metal  behaves  towards  water  and  oxygen  like  sodium  (p.  548). 
It  unites  directly  and  vigorously  with  hydrogen  (LiH),  nitrogen 
(Li8N),  and  oxygen  (Li20),  forming  stable  compounds.  The  chloride 
crystallizes  in  octahedra  (p.  549).  The  relative  insolubility  (p.  544) 
of  the  hydroxide  (LiOH),  the  carbonate  (LijC08),  and  the  phosphate 
(Li8P04,2H20),  is  in  sharp  contrast  to  the  easy  solubility  of  the  cor- 
responding compounds  of  the  other  alkali-metals,  and  links  lithium 
with  magnesium.  The  compounds  of  lithium  give  a  bright-red  color 
to  the  Bunsen  flame,  and  a  bright-red  and  a  somewhat  less  bright  orange 
line  are  seen  in  the  spectrum.  The  carbonate  is  used  in  medicine. 

IONIC  EQUILIBRIUM,  CONSIDERED  QUANTITATIVELY. 

The  Simple  Case.  —  In  view  of  the  predominance  of  ionic  actions 
in  the  chemistry  of  the  metals,  and  of  the  determinative  effect  of  ionic 
equilibria  on  many  actions,  it  is  essential  that  we  should  be  prepared 
in  future  for  a  more  exact  study  of  these  phenomena  than  we  have 
hitherto  attempted.  The  whole  basis  for  this  exact  study  has  already 
been  supplied,  and  only  more  specific  application  of  the  principles  is 
demanded. 

In  the  first  place,  the  principles  themselves  must  be  recalled. 
When  acetic  acid,  for  example,  is  dissolved  in  water,  it  is  ionized  thus  : 


The  amount  of  molecular  acetic  acid  dissociated  per  second  in  a  given 
amount  of  the  solution  is  proportional  to  the  concentration  of  the  mole- 
cules (  C^),  while  the  amount  of  the  two  ionic  materials,  hydrion  and 
acetanion,  uniting  to  form  molecules  of  acetic  acid  depends  on  the 
frequency  of  the  encounters  of  the  two  kinds  of  ions  and  is  proper- 


IONIC  EQUILIBRIUM,   CONSIDERED   QUANTITATIVELY       579 


tional  to  the  ionic  concentrations  (p.  297).  The  unit  of  concentration 
(p.  250)  is  1  mole  per  liter,  or,  in  the  present  case,  60  g.  of  the  acid, 
1  g.  of  hydrion,  and  59  g.  of  acetanion  respectively,  per  liter,  for  these 
numbers  represent  the  weight  of  one  mole  of  each  component.  Accord- 
ing to  the  law  of  concentration  (p.  252)  : 

<7,X  0, 


K, 


(1) 


and  the  numerical  value  of  this  fraction,  or  of  K,  remains  unchanged 
whatever  the  total  concentration  of  the  solution  may  be.  If  the  solution 
is  diluted,  for  example,  C2  and  Cz  diminish  relatively  less  quickly  than 
Cl  in  order  that  the  value  of  the  whole  expression  may  remain  the 
same.  This  is  accomplished  by  ionization  of  a  part  of  the  material 
whose  concentration  is  (7t  and  its  transference  to  the  ionic  forms  whose 
concentrations  are  C2  and  Ca,  respectively  (p.  297). 

A  numerical  example  will  show  that  this  law  of  concentration  ex- 
presses the  facts  with  considerable  exactness.  The  data  in  regard  to 
acetic  acid  are  as  follows  (p.  328)  : 


Acetic  Acid. 

I. 

Molar  Con- 
centra- 
tion. 

II. 

Proportion 
of  Whole  Mate- 
rial Ionized. 

III. 

Molar  Concentrations 
ofH'  (C2) 
and  of  C2H,02'  (C3). 
(I  X  II.) 

IV. 

Molar  Concentra- 
tion of 

TW 

Uni-molar     . 

1.0 

0.004 

0.004 

1.0  -  0.004 

Deci-molar    . 

0.1 

0.013 

0.0013 

0.1  -0.0013 

Centi-inolar  . 

0.01 

0.0407 

0.000407 

0.01  -0.000407 

Now  C  2  =  <73,  since  the  ions  are  produced  in  equal  numbers.  Also, 
for  our  purpose,  the  numbers  to  be  subtracted  in  column  iv  are  rela- 
tively so  small  that  the  values  1,  0.1,  and  0.01  may  be  taken  to  repre- 
sent Ci  without  appreciable  error.  Hence,  substituting  the  data  in 
equation  (1)  above,  we  have  : 

=  .  0,166. 


.i.  1 

In  other  words,  although  the  last  solution  is  a  hundred  times  more 
dilute  than  the  first,  and  the  degree  of  ionization  has  increased  ten 
times,  the  whole  expression  remains  close  to  the  value  0.04165  and  is 
essentially  constant. 

< 


580  INORGANIC  CHEMISTRY 

When  conductivity  data,  like  the  above,  are  applied  in  the  same 
way  to  the  cases  of  more  highly  ionized  substances,  the  values  of  K 
arc  less  nearly  constant.  It  is  supposed  that  with  this  class  of  sub- 
stances the  measurements  of  degrees  of  ionization  are  less  accurate, 
although  the  cause  of  the  discrepancy  has  not  been  fully  determined. 
However,  in  the  very  general  applications  of  the  data,  which  are  all 
that  we  shall  be  required  to  make,  the  conclusions  will  not  be  affected 
by  this  fact. 

The  expression  for  ionic  equilibrium  in  terms  of  concentration  is 
often  written  in  the  following  fashion  : 


[HC2H3OJ 


r, 


Here  [H*]  represents  the  molar  concentration  of  hydrion  ((72), 
[C2H30/]  that  of  acetanion  (<73),  and  [HC2H302]  that  of  molecular 
acetic  acid  (6\).  This  form  is  more  convenient  for  many  reasons,  and 
will  be  employed  frequently. 

Excess  of  One  Ion.  —  In  the  case  of  cupric  bromide  (p.  335),  we 
showed  that  increasing  the  concentration  of  the  bromidion  displaced 
the  equilibrium  by  favoring  the  union  of  the  ions  to  form  molecular 
cupric  bromide  :  2Br'  +  Cu"  —  >  CuBr2.  This  we  speak  of  as  a 
repression  of  the  ionization  of  the  cupric  bromide.  Now,  if  the  sub- 
stance is  a  slightly  ionized  one,  like  a  weak  acid  or  a  weak  base,  the 
repression  of  the  ionization  through  the  formation  of  molecules  in 
this  way  may  remove  so  many  of  that  one  of  the  ions  which  is  not 
present  in  excess  (corresponding  to  the  Cu"  in  the  foregoing  illus- 
tration), that  the  mixture  will  no  longer  respond  to  tests  for  the  ion 
so  removed.  This  is  an  interesting  and  very  common  case.  The 
behavior  of  acetic  acid,  a  weak,  slightly  ionized  acid,  will  serve  as  an 
illustration. 

In  normal  solution  (60  g.  in  1  1.)  acetic  acid  is  only  .004  ionized 
(p.  330),  so  that,  in  the  equation  for  the  equilibrium, 

(.996)  HC2H302<=?H'  (.004)  +  C2H3<V  (.004), 

the  relative  proportions  are  as  shown  by  the  numbers  in  parenthesis. 
If  the  whole  of  the  acid  (60  g.)  were  ionized,  there  would  be  1  g.  of 
hydrion  per  liter.  Yet,  even  in  this  much  smaller  concentration 
(.004  g.  per  liter),  the  acid  taste  of  the  H*  and  its  effect  upon  indi- 
cators can  be  distinctly  recognized.  If,  now,  solid  sodium  acetate  is 


IONIC   EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY        581 

dissolved  in  the  solution,  the  liquid  no  longer  gives  an  acid  reaction 
with  one  of  the  less  delicate  indicators,  like  methyl  orange  (q.v.). 
The  explanation  is  simple.  Sodium  acetate  is  highly  ionized.  It 
gives,  therefore,  a  large  concentration  of  acetanion  to  a  liquid  formerly 
containing  very  little.  This  causes  a  greatly  increased  union  of  the 
H*  ions  and  C2H302'  ions  and  the  former,  being  already  very  few  in 
number,  disappear  almost  entirely.  Hence  the  solution  becomes,  to 
all  intents  and  purposes,  neutral.  There  is  no  less  acetic  acid  present 
than  before,  but  the  concentration  of  hydrion  is  very  much  smaller. 

Formulation  and  Quantitative  Treatment  of  the  Case  of 
Excess  of  One  Ion.  —  If  the  semi-mathematical  mode  of  formulating 
an  equilibrium  (p.  580),  as  applied  to  the  case  of  an  ionogen,  be 
employed  here,  the  foregoing  general  statements  may  be  made  more 
precise  and  the  conclusions  clearer.  If  [H*]*and  [C2HjOa']  represent 
the  molecular  concentrations  of  hydrion  and  acetanion,  respectively, 
and  [HC2H302]  that  of  the  acetic  acid  molecules  at  equilibrium,  then  : 


X  _ 


[HC2H3OJ 

The  value  of  K  is  constant,  whether  the  strength  of  the  solution  of 
acetic  acid  is  great  or  small,  and  even  when  another  substance  with 
a  common  ion  is  present.  In  the  latter  case,  [C2H302']  and  [H'] 
stand  for  the  whole  concentrations  of  each  of  these  ionic  substances 
from  both  sources. 

Now,  as  we  have  seen,  in  normal  acetic  acid  [H*]  =  .004, 
[C2H30/]  =  .004  (for  the  number  of  each  kind  of  ions  is  the  same), 
and  [HC2H3OJ  =  .996,  practically  1.  Substituting  in  the  formula  : 

0.004  x  0.004  =jr(  =  fto<16). 

When  sodium  acetate  is  dissolved  in  the  liquid  until  the  solution  is 
normal  in  respect  to  this  substance  also,  the  following  additional 
equilibrium  has  to  be  considered  : 

(.47)  NaC2H302<=>Na'  (.53)  +  C2H302'  (.63). 

The  concentration  of  acetanion  from  this  source  is  .53,  so  that,  in  the 
mixture  of  acid  and  salt,  the  concentration  of  acetanion  [C2H30/]  will 
be  .53  +  .004  =.534,  or  nearly  134  times  larger  than  in  the  acid  alone. 
Hence,  in  order  that  the  product  [H*]  X  [C2H302']  may  recover,  as  it 


582  INORGANIC   CHEMISTRY 

must,  a  value  much  nearer  to  the  old  one,  [H*]  must  be  diminished  to 
something  like  T^?  of  its  former  magnitude.  That  is,  [H'J  will  become 
equal  to  about  0.00003,  the  rest  of  the  hydrion  uniting  with  a  corre- 
sponding amount  of  the  acetanion  to  form  molecular  acetic  acid.  The 
effect  of  adding  this  amount  of  sodium  acetate  therefore  is,  as  we 
have  seen,  to  reduce  the  concentration  of  the  hydrion  below  the 
amount  which  can  be  detected  by  use  of  an  indicator  like  methyl- 
orange. 

This  effect  is  of  course  reciprocal,  and  the  ionization  of  the  sodium 
acetate  will  be  reduced  also.  But  the  acetanion  furnished  by  the 
acetic  acid  is  relatively  so  small  in  amount  (.00003  against  .53)  that 
the  effect  it  produces  on  the  ionization  of  the  salt  is  imperceptible. 

It  will  be  noted  that  the  acetanion  and  hydrion  disappear  in 
equivalent  quantities,  for  they  unite.  There  is,  however,  so  much  of 
the  former  that  its  loss"  goes  unremarked,  while  there  is  so  little  of 
the  latter  that  almost  none  of  it  remains.  When  substances  of  more 
nearly  equal  degrees  of  ionization  are  used,  both  effects  are  equally  in- 
conspicuous. Thus,  sodium  chloride  and  hydrogen  chloride  in  normal 
solutions  yield  approximately  equal  concentrations  of  chloridion  (.784 
and  .676).  Hence,  if  one  mole  of  sodium  chloride  were  to  be  dissolved 
in  the  portion  of  water  already  containing  one  mole  of  hydrogen  chlo- 
ride, the  concentration  of  the  chloridion,  at  a  very  rough  estimate, 
would  be  nearly  doubled.  If  this  doubling  of  the  concentration  of 
chloridion  almost  halved  that  of  the  hydrion  (.784),  in  order  that  the 
expression  [Cl']  x  [H7]  -f-  [HC1]  might  remain  constant,  the  concentra- 
tion of  the  hydrion  would  still  be  about  .400  and  therefore  100  times 
as  great  as  in  molar  acetic  acid.  It  is  thus*  altogether  impossible 
to  reduce  the  concentration  of  the  hydrion  given  by  an  active  acid  like 
hydrochloric  acid  below  the  limit  at  which  indicators  are  affected,  for 
there  is  no  way  of  introducing  the  enormous  concentration  of  the  other 
ion  which  the  theory  demands. 

With  more  crude  means  of  observation  than  indicators  afford, 
effects  like  this  last  may  sometimes  be  rendered  visible.  This  was  the 
case  with  cupric  bromide  solution,  to  which  potassium  bromide  was 
added  (p.  335).  The  blue  of  the  dicuprion  disappeared  from  view, 
while  much  dicuprion  was  still  present,  because  the  brown  color  of  the 
molecular  cupric  bromide  covered  it  up  completely. 

Special  Case  of  Saturated  Solutions.  —  The  commonest  as 
well  as  the  most  interesting  application  of  the  conceptions  developed 


IONIC    EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY        583 

above  is  met  with  in  connection  with  saturated  solutions,  especially 
those  of  relatively  insoluble  substances. 

The  situation  in  a  system  consisting  of  the  saturated  solution  and 
excess  of  the  solute  has  been  discussed  already  (p  339).  In  the  case 
of  potassium  chlorate,  for  example,  we  have  the  following  scheme  of 
equilibria : 

KC103  (solid)  <=>  KC1O3  (diss'd)  ?=*  K'  +  CIO/. 

Solution  of  the  solid  is  prompted  by  the  solution  pressure  of  the  mole- 
cules, while  it  is  opposed  by  the  osmotic  pressure  of  the  dissolved  sub- 
stance, and  the  solution  is  saturated  when  these  tendencies  produce 
equal  effects  (p.  153).  Now  it  must  be  noted  that  the  tendency  di- 
rectly opposed  to  the  solution  pressure  is  the  partial  osmotic  pressure 
of  the  dissolved  molecules  alone.  The  chief  contents  of  the  solution, 
the  molecules  and  two  kinds  of  ions  of  the  salt,  and  any  foreign  mate- 
rial that  may  be  present,  are  like  a  mixture  of  gases,  and  the  principle 
of  partial  pressure  (p.  88)  is  to  be  applied.  The  ions  and  the  foreign 
material  do  not  deposit  themselves  upon  the  solid,  and  take,  therefore, 
no  part  directly  in  the  equilibrium  which  controls  solubility.  In 
respect  to  this  the  ions  are  themselves  foreign  substances.  Hence 
the  conclusion  may  be  stated  that,  in  solutions  saturated  at  a  given 
temperature  by  a  given  solute,  the  concentration  of  the  dissolved  mole- 
cules of  the  solute  considered  by  themselves  will  be  constant  what- 
ever other  substances  may  be  present. 

The  total  solubility  of  a  substance,  as  we  have  used  the  term 
hitherto,  is  made  up  of  a  molecular  and  an  ionic  part.  The  latter,  as 
we  shall  presently  see,  is  not  constant  when  a  foreign  substance  con- 
taining a  common  ion  is  already  in  the  liquid.  Since  the  treat- 
ment of  the  subject  requires  us  now  to  distinguish  between  the 
two  portions  of  the  solute,  a  diagram  (Fig.  101)  will  assist  in  empha- 
sizing the  distinction.  The  material  at  the  bottom  is  the  salt.  The 
molecules  and  ions  are  to  be  thought  of  as  being  mixed  and  as  being 
present  in  numbers  represented  by  the  factors  n  and  ra.  Since  no 
foreign  body  is  present,  the  two  ions  in  this  case  are  equal  in 
number. 

When  we  now  apply  these  ideas  to  the  mathematical  expression  of 
the  relation : 

[K-]  X  [CIO/] 
[KC10J 


584 


INORGANIC   CHEMISTRY 


we  perceive  that,  in  a  saturated  solution,  [KC103],  the  concentration 
of  the  molecules,  is  constant.     Transposing,  we  have 


X  [C103']  = 

Hence  the  relation  leads  to  the  important  conclusion  that,  in  a  satu- 
rated solution  the  product  of  the  molar  concentrations  of  the  ions  is 
constant.*  This  product  is  called  the  ion-product 
constant  for  the  substance.  The  law  of  the  con- 
stancy of  the  ion-product  in  a  saturated  solution 
is  one  of  the  most  useful  of  the  principles  of 
chemistry.  It  enables  us  to  explain  all  the  varied 
phenomena  of  precipitation  and  of  the  solu- 
tion of  precipitates  in  a  consistent  manner.  These 
applications  of  the  principle  will  be  explained 
in  the  next  chapter.  One  curious  kind  of  pre- 
cipitation will  be  described  here,  however,  as  an 
illustration  of  the  use  of  the  principle. 


»K'+mCI08' 


II 


n  KClOa 


11 


FIG.  101.  Illustration  of  the  Principle  of  Ion- fro- 

duct  Constancy.  —  When,  to  a  saturated  solution 
of  one  of  the  less  soluble  salts,  a  strong  solution  of  a  salt  having  one 
ion  in  common  with  the  first  salt  is  added,  precipitation  of  the  first 
salt  frequently  takes  place.  This  happens,  for  example,  with  a 
saturated  solution  of  potassium  chlorate,  which  is  not  very  soluble 
(molar  solubility  .52,  see  Table).  The  concentrations  [K*]  and  [CIO/] 
being  small,  one  may  easily  increase  the  value  for  one  of  the  ions,  say 
[CIO/],  fivefold,  by  adding  a  chlorate  which  is  sufficiently  soluble. 
To  preserve  the  value  of  the  product  [K']  X  [CIO/],  the  value  of  [K*] 
will  then  have  to  be  diminished  at  once  to  one-fifth  of  its  former 
value.  This  can  occur  only  by  union  of  the  ionic  material  it  repre- 
sents with  an  equivalent  amount  of  that  for  which  [CIO/]  stands.  The 
molecular  material  so  produced  will  thus  tend  at  first  to  swell  the 
value  of  [KC103].  But  the  value  of  [KC103]  cannot  be  increased,  for 
the  solution  is  already  saturated  with  molecules,  so  that  the  new 
supply  of  molecules,  or  others  in  equal  numbers,  •will  be  precipitated. 

*  The  principle  of  constant  concentration  of  dissolved  molecules,  stated 
above,  has  been  shown  to  express  the  facts  very  inaccurately.  Now  the  principle 
of  the  constancy  of  the  ratio  of  the  ion-product  to  the  concentration  of  the  mole- 
cules is  also  inaccurate,  yet  in  such  a  way  that  the  two  errors  neutralize  one 
another.  Thus,  the  principle  of  ion-product  constancy  here  given  is  in  itself  fairly 
exact. 


IONIC   EQUILIBRIUM,    CONSIDERED   QUANTITATIVELY         585 

Hence  the  ionic  part  of  the  dissolved  substance  may  be  diminished, 
the  equilibria  (p,  583)  may  be  partially  reversed,  and  we  may  actually 
precipitate  a  part  of  the  dissolved  material  without  introducing  any 
substance,  which,  in  the  ordinary  sense,  can  interact  with  it. 

In  point  of  fact,  when,  to  a  saturated  solution  of  potassium  chlo- 
rate there  is  added  a  saturated  solution  of  potassium  chloride  (KC1)  or 
of  sodium  chlorate  (NaC103),  a  precipitate  of  potassium  chlorate  is 
thrown  down.  These  two  salts,  containing  each  one  of  the  ions  of 
KC103,  and  being  much  more  soluble  than  the  latter  (see  Table), 
increase  the  concentration  of  one  ion  and  cause  the  precipitation  in 
the  fashion  just  explained. 

The  product  of  the  concentrations  of  the  ions,  for  example  [K']  X 
[CIO/],  is  called  also  the  solubility  product,  because  these  two  values 
jointly  determine  the  magnitude  of  the  solubility  of  the  substance. 
The  solubility  of  the  molecules  is  irreducible,  but  the  ionic  part  of  the 
dissolved  material  may  become  vanishingly  small  if  the  value  of  either 
[X'J  or  [¥']  is  very  minute.  The  ionic  part  of  any  particular  sub- 
stance is  made  up  of  the  smaller  of  the  two  concentrations  of  the 
ionic  substances  which  it  yields,  plus  an  equivalent  amount,  and  no 
more,  of  the  concentration  of  the  other  ion.  The  rest  of  the  other 
ionic  substance  is  part  of  the  solubility  of  some  other  component. 

Other  Illustrations.  —  The  precipitation  of  sodium  chloride 
from  a  saturated  solution,  by  the  introduction  of  gaseous  hydrogen 
chloride  (p.  570),  is  to  be  explained  in  the  same  manner.  The 
equilibria : 

Nad  (solid)  <=>  NaCl  (diss'd)  ^±  Na"  +  Cl' 

are  reversed  by  the  introduction  of  additional  Cl'  from  the  very 
soluble,  and  highly  ionized  HOI. 

The  evolution  of  a  steady  stream  of  hydrogen  chloride  is  often 
accomplished  by  allowing  concentrated  sulphuric  acid  to  flow  into 
saturated  hydrochloric  acid : 

H-  +  Cl'  <=>  HC1  (diss'd)  <±  HC1  (gas). 

The  effect  is  due  in  part  to  repression  of  the  ionization  of  the  hydro- 
gen chloride  and  elimination  of  molecules  of  the  gas  from  the  water 
which  is  already  saturated  with  molecules  of  the  same  kind.  The 
"  salting  out "  of  soap  (p.  505)  is  probably  a  similar  phenomenon. 

Exercises.  —  1.  The  vapor  density  of  sodium  peroxide  has  not 
been  determined.  Why  is  the  formula  Na2O2  assigned  to  it  ? 


586  INORGANIC   CHEMISTRY 

2  Construct  a  scheme  of  equilibria  (p.  371)  showing  the  hydroly- 
sis of  calcium  sulphide.  Why  does  the  presence  of  calcium  hydroxide 
diminish  the  tendency  to  hydrolysis  ? 

3.  Show  the  application  of  the  principle  of  ion-product  constancy 
to  the  salting  out  of  soap  (p.  505). 

4.  What  will  be  the  effect  of  adding  a  concentrated  solution  of 
silver  nitrate  to  a  saturated  solution  of  silver  sulphate  or  of  silver 
acetate  (see  Table  of  solubilities)  ? 


CHAPTER  XXXV 
THE  METALS   OF  THE   ALKALINE   EARTHS 

The  Chemical  Relations  of  the  Elements.  —  The  metals  of  this 
group,  calcium  (Ca,  at.  wt.  40.1),  strontium  (Sr,  at.  wt.  87.6),  and  ba- 
rium (Ba,  at.  wt.  137.4),  constitute  a  typical  chemical  family  both  in  the 
qualitative  resemblance  to  one  another  of  the  elements  and  correspond- 
ing compounds  and  in  the  quantitative  variation  in  the  properties  with 
increasing  atomic  weight.  The  metals  themselves  displace  hydrogen 
vigorously  from  cold  water,  giving  hydroxides.  The  solutions  of  these 
hydroxides,  although  dilute,  on  account  of  a  rather  small  solubility, 
are  strongly  alkaline  in  reaction.  The  high  degree  of  ionization  of  the 
hydroxides  recalls  the  hydroxides  of  the  metals  of  the  alkalies,  and 
their  relative  insolubility  the  hydroxides  of  the  "earths"  (q.v.). 

In  all  their  compounds,  calcium,  strontium,  and  barium  are  bivalent. 
The  hydroxides  are  formed  by  union  of  the  oxides  with  water,  and,  ex- 
cept in  the  case  of  barium  hydroxide,  are  easily  decomposed  again  by 
heating.  The  carbonates,  when  heated,  yield  the  oxide  of  the  metal 
and  carbon  dioxide,  barium  carbonate  being  the  most  difficult  to  decom- 
pose. The  nitrates  are  broken  up  by  heating  and  yield  the  oxide  of  the 
metal,  nitrogen  tetroxide,  and  oxygen.  In  these  and  other  respects 
the  compounds  of  the  metals  of  the  alkaline  earths  resemble  those  of 
the  heavy  metals  and  differ  from  those  of  the  metals  of  the  alkalies. 
Barium  approaches  the  latter  most  nearly. 

The  table  on  p.  544  shows  that  the  chlorides  and  nitrates  of  calcium, 
strontium,  and  barium  are  all  soluble  in  water,  the  solubility  dimin- 
ishing in  the  order  given.  The  sulphates  and  hydroxides  cover  a  wide 
range  from  slight  solubility  to  extreme  insolubility.  Of  the  sulphates, 
2000,  110,  and  2.3  parts,  respectively,  dissolve  in  one  million  parts  of 
water.  In  the  case  of  the  hydroxides  the  order  of  magnitude  is  re- 
versed, and  the  corresponding  numbers  are  1700,  7700,  and  37,000. 
The  carbonates  are  almost  as  insoluble  as  is  barium  sulphate.  The 
new  element,  radium  (Ra,  at.  wt.  226.5),  appears  to  belong  to  this 
family  (see  under  Uranium). 

587 


588  INORGANIC   CHEMISTRY 

CALCIUM. 

Occurrence.  —  The  fluoride,  and  the  various  forms  of  the  carbon 
ate,  sulphate,  and  phosphate  which  are  found  in  nature,  are  described 
below.  As  silicate,  calcium  occurs,  along  with  other  metals,  in  many 
minerals  and  rocks.  It  is  found  also  in  plants,  and  its  compounds 
are  important  constituents  of  the  bones  and  shells  of  animals. 

The  Metal. — Although  the  alkali  metals  can  be  liberated  by  heating 
the  carbonates  with  carbon,  the  metals  of  the  present  family  are  not 
obtainable  by  this  means.  This  may  be  due,  in  part,  to  imperfect  con- 
tact between  the  materials  in  consequence  of  the  infusibility  of  the  ox- 
ides. Calcium  is  most  easily  made  by  electrolysis  of  the  molten 
chloride.  A  hollow  cylinder  made  of  blocks  of  carbon  bolted  together 
and  open  above,  forms  the  anode.  A  rod  of  copper  hanging  so  that  its 
end  dips  into  the  melt  forms  the  cathode.  The  melting  of  the  anhy- 
drous calcium  chloride  with  which  the  cylinder  is  filled  is  started  by 
means  of  a  thin  rod  of  carbon  laid  across  from  the  anode  to  the  cath- 
ode. When  the  heat  generated  by  the  passage  of  the  current  through 
this  highly  resisting  medium  has  melted  a  sufficient  amount  of  the  salt, 
the  rod  is  removed,  and  the  resistance  of  the  fused  material  suffices  to 
maintain  the  temperature.  The  calcium  rises  round  the  cathode  and 
collects  on  the  surface  of  the  bath.  By  slowly  elevating  the  copper 
cathode,  the  calcium,  which  adheres  to  it,  may  be  drawn  out  of  the 
fused  mass  in  the  form  of  a  gradually  lengthening,  irregular  rod.  The 
rod  of  calcium  is  kept  constantly  in  contact  with  the  metal  which 
accumulates  on  the  surface,  and  thus  forms  one  of  the  electrodes. 

Calcium  is  a  silver-white,  crystalline  metal  (m.-p.  760°,  sp.  gr.  1.85) 
which  is  a  little  harder  than  lead,  and  can  be  cut,  drawn,  and  rolled. 
It  interacts  violently  with  water.  When  dry  and  cold  it  is  inactive, 
but  when  heated  it  unites  vigorously  with  hydrogen,  oxygen,  the  halo- 
gens, and  nitrogen.  It  burns  in  the  air,  giving  a  mixture  of  the  oxide 
and  nitride.  The  presence  of  the  latter  may  be  shown  by  the  libera- 
tion of  ammonia  when  water  is  brought  in  contact  with  the  residue : 

6H2O  ->  3Ca(OH)2 


Calcium  Hydride.  —  The  hydride  CaH2,  when  formed  by  direct 
union  of  the  constituents,  is  a  white  crystalline  body.  It  interacts 
vigorously  with  water,  liberating  hydrogen. 


THE  METALS  OF  THE  ALKALINE  EARTHS        589 

Calcium  Chloride.  —  Calcion  is  present  in  small  amount  in 
sea- water,  and  hence  compounds  containing  calcium  chloride,  such  as 
tachydrite  CaCl2,MgCl2,12H20,  are  found  in  salt  deposits.  The  salt, 
for  which  there  is  no  extensive  commercial  application,  is  formed  in 
large  quantities  as  a  by-product  in  several  industrial  operations.  Thus, 
it  arises  in  the  liberatioD  of  ammonia  from  ammonium  chloride  by  the 
action  of  lime,  in  the  manufacture  of  potassium  chlorate  (p.  273),  and 
in  the  Solvay  soda  process  (p.  574).  By  evaporation  of  any  solution  the 
hexahydrate  of  the  salt,  CaCl2, 6H20,  is  obtained  in  large,  deliquescent, 
six-sided  prisms.  On  account  of  the  great  concentration  of  a  satu- 
rated solution  of  this  compound,  the  solid  and  solution  do  not  reach  a 
condition  of  equilibrium  with  ice  (cf.  p.  164)  until  the  temperature  has 
fallen  to  —48°.  The  freezing  mixture  must  be  made  with  the  hydrate, 
and  not  with  the  anhydrous  salt,  as  the  latter  gives  out  much  heat  in 
dissolving.  The  former,  on  the  other  hand,  absorbs  heat  in  liquefying, 
as  all  solids  do. 

There  are  several  other  hydrates  of  calcium  chloride  containing 
less  than  6H2O,  and  those  containing  less  water  have  lower  aqueous 
tensions  (cf.  p.  122)  than  those  containing  more.  By  elevating  the  tem- 
perature, however,  it  is  easy  to  raise  the  aqueous  tension  even  of  the 
monohydrate  until  it  exceeds  the  partial  pressure  of  water  vapor  in 
ordinary  moist  air,  and  so  to  drive  out  the  water.  To  perform  this 
rapidly,  a  temperature  of  over  200°  is  required.  The  dehydrated 
calcium  chloride  forms  a  porous  mass  which  is  used  in  chemical  labo- 
ratories for  drying  gases  and  liquids.  Usually  the  dehydration  is  left 
incomplete,  as,  at  the  temperature  required  to  complete  it  rapidly,  some 
interaction  with  the  water  occurs  (CaCl2  +  H20  — >  CaO  -f-  2HC1),  and 
a  little  free  alkali  is  present  in  the  product.  When  calcium  chloride 
is  used  as  a  drying  agent,  it  is  naturally  able  to  reduce  the  partial 
pressure  of  the  water -vapor  only  to  the  value  of  the  aqueous  tension  of 
the  hydrate  which  is  present,  and  no  further.  Even  at  low  tempera- 
tures the  aqueous  tensions  of  hydrates  are  always  perceptible  (cf. 
p.  121).  Concentrated  sulphuric  acid  is  a  more  thorough  drying  agent 
than  calcium  chloride,  and  phosphorus  pentoxide,  whose  hydrated  form 
(metaphosphoric  acid)  has  no  observable  aqueous  tension,  is  better  still. 

Calcium  chloride  forms  molecular  compounds,  not  only  with  water, 
but  also  with  ammonia  (CaCl2,8NH8)  and  with  alcohol.  For  drying 
these  substances,  therefore,  quicklime  is  employed.  Hydrogen  sulphide 
interacts  with  the  salt,  giving  hydrogen  chloride,  which  renders  the 
gas  impure.  This  gas  is  therefore  dried  with  phosphorus  pentoxide. 


590  INORGANIC   CHEMISTRY 

Calcium  Fluoride.  —  This  compound  occurs  in  nature  as  fluorite 
or  fluor-spar  CaF2.  It  crystallizes  in  cubes,  is  insoluble  in  water,  and 
when  pure  is  colorless.  Natural  specimens  often  possess  a  green  tint 
or  show  a  violet  fluorescence.  It  is  formed  as  a  precipitate  when  a 
soluble  fluoride  is  added  to  a  solution  of  a  salt  of  calcium. 

Fluorite  is  used  in  the  etching  of  glass,  as  the  source  of  the  hydrogen 
fluoride  (p.  241).  It  is  easily  fusible,  as  its  name  indicates  (L&t.ftuere, 
to  flow),  and  is  employed  in  metallurgical  operations,  for  the  purpose 
of  lowering  the  melting-point  (cf.  p.  163)  of  the  slag  (q.v.),  and  so  facili- 
tating the  separation  of  the  latter  from  the  metal. 

Calcium  Carbonate.  —  This  compound  is  found  very  plentifully  in 
nature.  Limestone  is  a  compact,  indistinctly  crystalline  variety,  while 
marble  is  a  distinctly  crystalline  form.  Chalk*  is  a  deposit  consist- 
ing of  the  calcareous  parts  of  minute  organisms  ;  and  egg-shells,  oyster- 
shells,  coral,  and  pearls  are  other  varieties  of  organic  origin.f  A 
laminated  kind  of  limestone  found  at  Solnhofen  is  used  for  lithographic 
work.  Calcite  and  Iceland  spar  (Ger.  spalten,  to  split)  are  pure  crys- 
tallized calcium  carbonate.  The  former  occurs  in  flat  rhombohedrons, 
or  in  pointed,  six-sided  crystals  (Fig.  52,  p.  138)  known  as  scalenohe- 
drons  ("dog-tooth"  spar)  belonging  to  the  same  system.  All  the 
crystals  split  with  ease  parallel  to  three  planes  of  cleavage,  giving  rhom- 
bohedrons of  the  shape  shown  in  Fig.  9  (p.  14),  but  this  nearly  cubical 
form  is  itself  seldom  found  in  nature.  An  entirely  different  crystal- 
lized variety  is  known  as  aragonite.  This  belongs  to  the  rhombic 
system,  although  complex  crystals  ("  twins  ")  of  hexagonal  outline  con- 
stitute the  most  familiar  specimens.  Aragonite,  when  heated  strongly, 
resolves  itself  into  a  mass  of  minute  crystals  of  calcite,  and  the  latter 
is  the  more  stable  form  of  the  substance.  When  calcium  carbonate  is 
produced  by  precipitation  it  is  at  first  amorphous  but  slowly  becomes 
crystalline.  In  cold  liquids  the  resulting  crystals  are  calcite  ;  but  in 
warm  solutions  the  less  stable  form,  that  of  aragonite,  is  first  assumed. 

When  heated,  calcium  carbonate  dissociates,  giving  carbon  dioxide 
and  quicklime : 

CaC08  <±  CaO  +  CO2. 

At  ordinary  temperatures  the  decomposition  is  imperceptible.     On  the 
contrary,  atmospheric  carbon  dioxide,  in  spite  of  its  very  low  partial 

*  Blackboard  "  crayon  "  is  usually  made  of  gypsum  and  not  of  chalk. 
t  The  hard  coverings  of  Crustacea  and  insects  are  not  made  of  this  substance, 
but  of  an  organic  material  called  chitin. 


THE  METALS  OF  THE  ALKALINE  EARTHS        591 

pressure,  combines  with  quicklime,  giving  "  air-slaked "  lime.  As 
the  temperature  rises,  however,  the  tension  of  carbon  dioxide  coming 
from  the  carbonate  increases,  and  has  a  fixed  value  for  each  temperature. 
If  it  is  continually  allowed  to  escape,  so  that  the  maximum  pressure  is 
not  reached,  the  whole  of  the  salt  eventually  decomposes.  If,  on  the 
other  hand,  the  gas  is  confined,  the  system  reaches  a  condition  of  equi- 
librium. Attempts  to  increase  the  pressure  on  the  gas  beyond  the 
dissociation  pressure  proper  to  the  existing  temperature,  result  in 
recombination.  The  phenomenon  is  precisely  similar  to  the  dissocia- 
tion of  barium  dioxide  (p.  257)  and  to  the  evaporation  of  a  liquid 
(p.  117).  The  following  data  show  the  actual  tensions  at  various 
temperatures  : 

Temperature  547°         625°         745°         810°         812°  865° 

Tension  in  mm.          27  56  289  678  753  1333 

Thus  at  812°  the  pressure  almost  reaches  one  atmosphere,  and  at  865° 
it  approaches  two  atmospheres. 

Limestone  is  used  in  the  manufacture  of  quicklime  (q.v.)  and  of 
glass.  It  is  employed  largely  as  a  flux  in  metallurgy,  when  minerals 
rich  in  silica  are  brought  into  fusible  form  by  the  production  of  calcium 
silicate  (CaSi03).  Large  amounts  also  find  application  as  building-stone. 

The  Phase  Rule,  a  Method  of  Classifying  all  Systems  in 
Equilibrium.  —  The  formal  resemblance  that  we  have  just  shown  to 
exist  between  the  modes  of  behavior  of  a  system  composed  of  water 
and  water-vapor  in  physical  equilibrium,  on  the  one  hand,  and  of  a 
system  made  up  of  calcium  oxide,  carbon  dioxide,  and  undecomposed 
calcium  carbonate  in  chemical  equilibrium,  on  the  other,  is  not  a  co- 
incidence. A  study  of  all  kinds  of  systems  in  equilibrium  shows  that 
their  different  modes  of  behavior  are  limited  in  variety  and  can  be 
classified  in  a  very  simple  way. 

The  categories  used  for  classification  are :  (1)  the  independent 
components  in  the  system,  and  their  number;  (2)  the  distinct, 
physically  separable  parts  or  phases  (p.  156)  of  the  system,  and  their 
number ;  and  (3)  the  conditions  —  temperature,  pressure,  and  concen- 
tration (or  volume)  —  and  the  degree  of  variability  in  the  conditions 
which  is  possible  without  the  occurrence  of  a  change  in  the  number  of 
the  phases. 

The  mode  of  employment  of  these  three  categories  may  be  illus- 
trated in  the  order  of  their  mention  : 


592  INORGANIC   CHEMISTRY 

1.  In  the  water  and  water-vapor  system,  water  is  the   only  com- 
ponent.    In  the  calcium  carbonate  system,  the  independent  components 
are  two  in  number,  calcium  oxide  and  carbon  dioxide. 

2.  In  the  water  and  water-vapor  system  there  are  two  phases,  the 
liquid  phase  and  the  vapor  phase.     In  the  calcium  carbonate  system 
there  are  three  phases  —  two  solid  phases,  the  carbonate  and  oxide,  and 
one  gaseous  phase,  the  carbon  dioxide. 

3.  In  the  water  and  water-vapor  system  either  the  temperature  or 
the  pressure  may   be  altered,  within  certain  limits,  at   will.      But, 
whichever  one  of  these  two  conditions  it  be  that  is  thus  changed,  the 
preservation  of  the  two  phases  will  at  once  require  a   simultaneous 
modification  in  the  other  condition,  of  such  a  nature  as  will  suit  the 
new  value  of  the  first.     Thus,  if  the  pressure  upon  the  vapor  is  raised, 
the  vapor  phase  will  be  destroyed  (p.  117)  unless  the  temperature  is 
simultaneously  elevated   to  a   certain  definite  point  (p.  116).     Sim- 
ilarly, if  the  temperature  is  raised,  the  liquid  phase  will  all  pass  into 
vapor  unless  a  sufficient  increase  in  the   pressure  is  simultaneously 
effected.     There  is  therefore  one,  and  only  one  degree  of  variability  in 
the    conditions  —  the  system  is  univariant.     By  a  study  of  the  cal- 
cium carbonate  system,  as  described  above,  it  will  be  seen  that  it  also 
is  a  univariant  system. 

A  partial  generalization  of  these  results  leads  to  the  conclusion 
that  when  the  number  of  the  phases  exceeds  the  number  of  the  compo- 
nents by  one,  the  system  is  univariant.  Additional  illustrations  are 
now  required  for  reaching  a  still  more  general  statement. 

If  ice  be  added  to  the  water  and  water-vapor  system,  and  the  system 
be  allowed  to  reach  equilibrium  with  all  three  phases  present,  we  find 
on  analyzing  as  before:  one  component  (water),  three  phases  •  (solid, 
liquid,  and  gaseous),  and  no  variability  in  the  system.  Neither  tem- 
perature nor  pressure  may  be  altered  without  ensuing  disappearance 
of  one  or  other  of  the  three  phases.  This  system  is  therefore 
invariant. 

It  thus  appears  that  with  an  equal  number  of  components,  the 
more  phases  we  have,  the  more  restricted  are  the  possibilities  of 
change  in  the  conditions.  When  the  number  of  phases  equals  the 
number  of  components,  the  system  is  bivariant ;  when  the  number  of 
phases  exceeds  the  number  of  components  by  one,  the  system  is  uni- 
variant ;  when  the  number  of  phases  exceeds  the  number  of  compo- 
nents by  two,  the  system  is  invariant,  and,  in  general, 

Phases  +  Variable  conditions  =  Components  +  2. 


THE  METALS  OF  THE  ALKALINE  EARTHS        593 

The  law,  of  which  this  equation  is  the  most  compact  expression,  is 
known  as  the  phase  rule,  and  was  first  formulated  by  Willard  Gibbs,  of 
Yale  University.  It  applies  to  physical  and  chemical  equilibria 
without  distinction,  and  involves  no  consideration  of  molecular  or  other 
hypotheses. 

Thus,  the  formal  resemblance  between  the  dissociation  phenomena 
exhibited  by  calcium  carbonate  and  other  compounds,  on  the  one  hand, 
and  the  behavior  of  a  liquid  in  contact  with  its  vapor  on  the  other,  is 
due  simply  to  the  fact  that  in  each  case  the  number  of  phases  exceeds 
the  number  of  components  by  one.  This  will  be  found  to  hold  in  all 
cases  where  there  is  at  each  temperature  a  constant  dissociation  press- 
ure. A  decomposing  hydrate,  for  example,  furnishes  such  a  case. 
The  system  is  made  up  of  one  gaseous  phase  (water-vapor)  and  two 
solid  phases  (the  hydrate,  and  the  anhydrous  substance  or  a  lower 
hydrate).  It  has  three  phases  and  two  components  (water  and  the 
anhydrous  substance),  and  is,  therefore,  univariant. 

Again,  when  we  have  a  sharp  transition-point  at  a  fixed  temper- 
ature, that  is,  a  unique  temperature  at  which  alone  several  different 
states  of  aggregation  of  a  substance  can  co-exist  (p.  115),  the  system 
is  always  invariant.  Thus,  ice  and  water  (and  vapor)  co-exist  at  the 
melting-point  of  ice:  three  phases  and  one  component.  Again,  ice, 
solid  sodium  chloride,  salt  solution,  and  water  vapor  all  co-exist  at 
—  22°,  the  transition-point :  four  phases  and  two  components  in  this 
case.  Still,  again,  at  96°  two  solid  forms  of  sulphur  (p.  368)  co-exist 
with  sulphur  vapor.  In  these  cases  the  change  which  takes  place  at  the 
transition  point  is  purely  physical.  Analogous  cases  in  which  the 
change  is  a  chemical  one  are  equally  familiar.  The  decahydrate  of 
sodium  carbonate  decomposes  (p.  575)  above  35.2°.  At  this  temper- 
ature the  decahydrate  and  the  monohydrate  co-exist  with  the  saturated 
solution  and  water-vapor  :  four  phases  and  two  components.  The  sys- 
tem is,  therefore,  invariant.  The  cases  of  gypsum  (see  p.  603)  and 
sodium  sulphate  (p.  576)  are  similar. 

Hard  Water.  —  Calcium  carbonate  is  almost  insoluble  in  water. 
In  the  cold  the  solubility  is  a  little  over  1  part  in  100,000 ;  in  hot 
water  the  solubility  is  even  smaller.  Water  containing  carbonic  acid 
dissolves  it  more  freely,  on  account  of  the  formation  of  the  more 
soluble  calcium  bicarbonate  (cf.  p.  482)  : 

CaC08  +  H2C08  <=>  Ca(HC08)2. 


594  INORGANIC    CHEMISTRY 

A  considerable  excess  of  carbonic  acid  is  required,  as  the  action  is 
markedly  reversible.  At  15°  a  liter  of  water  saturated  with  carbon 
dioxide  (at  760  mm.  pressure)  dissolves  0.385  g.  of  the  carbonate, 
whereas  a  liter  of  pure  water  dissolves  but  .013  g.  With  a  higher 
pressure  of  carbon  dioxide  still  larger  amounts  may  be  dissolved.  Con- 
versely, when  the  carbon  dioxide  is  driven  out  by  boiling,  the  carbon- 
ate is  reprecipitated. 

Water  containing  salts  of  lime  or  magnesia  in  solution  is  known  as 
hard  water.  The  carbonate,  which  may  be  thrown  down  by  boiling, 
gives  "  temporary  hardness,"  while  the  sulphate  of  calcium,  since  it  is 
soluble  per  se  and  is  not  affected  by  boiling,  gives  "  permanent  hard- 
ness." Soap  will  not  produce  a  lather  with  hard  water  until  the 
calcion  has  all  been  precipitated  in  the  form  of  the  calcium  salts  of 
stearic,  palmitic,  and  oleic  acids  (p.  506).  Commercially,  the  hardness 
is  therefore  measured  by  the  quantity  of  a  standard  soap  solution 
which  is  just  sufficient  to  produce  a  persistent  lather  in  a  given  amount 
of  the  water. 

The  temporary  hardness  may  be  removed  by  adding  to  the  water  a 
quantity  of  slaked  lime  sufficient  to  convert  the  excess  of  carbonic  acid 
into  calcium  carbonate.  The  permanent  hardness  may  be  destroyed  by 
the  addition  of  sodium  carbonate.  In  both  cases  the  precipitated  car- 
bonate is  allowed  to  settle  or  is  separated  by  filtration. 

When  evaporated  in  steam-boilers  hard  water  leaves  behind  a  heavy 
deposit  of  boiler-crust  containing  all  the  salts  formerly  in  solution. 
Similar  water,  when  it  dries  slowly  on  the  roofs  of  caverns,  gives  rise  to 
stalactites.  The  drippings  form  stalagmites  on  the  floors. 

Calcium  Oxide.  —  Pure  oxide  of  calcium  may  be  made  by  ignition 
of  pure  marble  or  calcite.  For  commercial  purposes  limestone  is  con- 
verted into  quicklime  in  kilns.  In  the  United  States  the  "  long-flame  " 
process,  in  which  the  kiln  is  first  charged  with  limestone  and  a  fire  is 
then  kindled  in  a  cavity  left  at  the  bottom,  is  the  one  most  commonly 
used.  Elsewhere,  the  limestone  and  coal  are  thrown,  in  alternate 
layers,  into  the  kiln,  and  the  products  are  withdrawn  at  the  bottom.  The 
latter,  the  "  short-flame  "  method,  demands  less  fuel,  since  the  operation 
is  continuous,  and  the  structure  is  never  allowed  to  cool,  but  the  quick- 
lime is  mixed  with  the  ash  of  the  coal.  In  both  cases  a  free  passage  of 
excess  of  air  through  the  mass  is  desirable  in  order  that,  by  reduction 
of  the  partial  pressure  of  the  carbon  dioxide,  the  reverse  action  (p.  257) 
may  be  minimized.  The  better  the  ventilation,  the  lower  the  tempera* 


THE  METALS  OF  THE  ALKALINE  EARTHS        595 

ture  at  which  rapid  decomposition  can  be  effected  (cf.  p.  592).  The 
lowest  temperature  that  will  suffice  is  employed,  as  strong  heating  need- 
lessly favors  the  interaction  of  the  calcium  oxide  with  the  clay  which 
most  limestone  contains.  The  product  of  the  interaction  with  clay, 
calcium  silicate,  is  fusible,  and  by  closing  many  of  the  pores  of  the 
quicklime  renders  the  subsequent  slaking  imperfect. 

Pure  calcium  oxide  is  a  white,  porous  solid.  It  is  infusible  even  in 
the  oxyhydrogen  flame,  but  may  be  melted  and  boiled  in  the  electric 
arc.  It  is  not  reducible  by  sodium,  or  by  carbon  excepting  at  the 
temperature  of  the  electric  furnace. 

Calcium  Hydroxide.  —  When  water  is  poured  upon  quicklime 
it  is  first  absorbed  into  the  pores  mechanically.  The  chemical  union 
by  which  the  hydroxide  is  formed  : 

CaO  +  H20  <±  Ca(OH)2, 

proceeds  slowly.  When  it  is  complete  the  product  is  a  fine  powder 
occupying  a  much  larger  volume  than  the  original  materials.  The 
action  is  accompanied  by  the  development  of  much  heat,  and  during 
its  progress  a  part  of  the  water  is  driven  off  as  steam.  The  action  is 
reversible,  and  at  a  high  temperature  the  hydroxide  can  be  dehy- 
drated. Quicklime  from  pure  limestone  slakes  easily,  and  is  known  as 
"  fat "  lime.  That  made  from  material  containing  clay  or  magnesium 
carbonate  is  "  poor  "  lime.  The  latter  slakes  slowly  and  often  incom- 
pletely, and,  when  used  for  mortar,  does  not  harden  so  satisfactorily. 

Calcium  hydroxide  is  slightly  soluble  in  water,  600  parts  of  water 
dissolving  1  part  of  the  hydroxide  at  18°,  and  about  twice  as  much 
water  being  required  at  100°.  The  solution,  relatively  to  its  concen- 
tration, is  strongly  alkaline.  On  account  of  its  cheapness,  this  sub- 
stance is  used  by  manufacturers  in  almost  all  operations  requiring  a 
base,  and  it  thus  occupies  the  same  position  amongst  bases  that  sul- 
phuric acid  does  amongst  acids.  When  the  presence  of  much  water  is 
unnecessary  or  undesirable,  a  suspension  of  the  solid  hydroxide  in  the 
saturated  solution  ("milk  of  lime"),  or  even  a  paste,  is  employed. 
In  such  cases,  as  in  the  manufacture  of  caustic  alkalies  (p.  585),  the 
action  takes  place  with  the  part  which  is  at  the  moment  in  solution, 
and  proceeds  through  the  continual  readjustment  of  a  complex  set  of 
equilibria.  Some  of  the  industries  in  which  caustic  lime  plays  a  part 
are  :  the  manufacture  of  alkalies,  bleaching  powder,  and  mortar,  the 
removal  of  the  hair  from  hides  in  preparation  for  tanning,  and  the 
purification  of  illuminating-gas  (p.  513). 


596  INORGANIC   CHEMISTRY 

Mortar  and  Cement.  —  Mortar  is  made  by  mixing  water  with 
slaked  lime  and  a  large  proportion  of  sand.  The  "  hardening  "  process 
consists  in  an  interaction  of  the  carbon  dioxide  of  the  air  with  tl.v 
calcium  hydroxide  : 

C02  +  Ca(OH)2  ->  CaC03  +  H2O. 

After  the  superficial  parts  have  been  changed,  the  process  goes  on 
very  slowly,  and  many  years  are  required  before  the  deeper  layers  have 
been  transformed.  The  minute  crystals  of  calcium  carbonate  which 
are  formed  are  interlaced  with  the  sand  particles,  and  a  rigid  mass  if1 
finally  produced.  The  sand  is  useful  in  two  ways.  In  the  first  place : 
it  makes  the  whole  material  more  porous,  and  so  facilitates  the 
diffusion  of  the  gas  into  the  interior.  In  the  second  place,  since  the 
sand  is  not  itself  altered,  its  presence  prevents  the  formation  of  large 
cracks  which  would  otherwise  arise  from  the  shrinkage  that  accom- 
panies the  formation  of  the  carbonate.  The  "  hardening "  does  not 
begin  until  the  excess  of  water  used  in  making  the  mortar  has  evapo- 
rated, and  hence  ordinary  mortar  is  unsuitable  for  use  in  damp  places 
such  as  cellars. 

Cement  is  made  by  strongly  heating  a  mixture  of  limestone,  clay, 
and  sand,  and  pulverizing  the  product.  Some  natural  limestones,  con- 
taining over  20  per  cent  of  clay,  give  cements  without  the  addition  of 
other  ingredients.  When  the  cement  is  mixed  with  water,  it  gradually 
sets  to  a  solid  mass  which  appears  to  consist  of  a  mixture  of  silicates 
of  calcium  and  aluminium.  The  change  proceeds  throughout  the 
whole  material  simultaneously,  since  it  is  not  dependent  on  access  of 
any  gas,  and  not,  as  in  the  case  of  mortar,  from  the  surface  inwards. 
For  this  reason  the  hardening  of  cement  occurs  just  as  well  under 
water  as  in  any  other  locality. 

Calcium  Oxalate. — This  salt  may  be  observed  under  the  micro- 
scope in  the  cells  of  many  plants.  It  appears  in  the  form  of  needle- 
shaped  or  of  granular  crystals.  Since  it  is  the  least  soluble  salt  of  cal- 
cium, its  formation  is  used  as  a  test  for  calcium  ions.  Calcium  is  es- 
timated quantitatively  by  adding  ammonium  oxalate  to  the  neutral  or 
slightly  alkaline  solution  of  the  calcium  salt.  The  precipitate  is  sep- 
arated by  filtration,  washed  with  water,  and  then  heated  strongly  (ignited) 
in  a  crucible.  The  product  weighed  is  calcium  oxide,  CaC204— >CaO 
-f  C02  -f  CO.  More  often,  perhaps,  the  oxalate  is  ignited  with  sul- 
phuric acid,  and  the  calcium  weighed  as  sulphate. 


THE  METALS  OF  THE  ALKALINE  EARTHS        597 

Theory  of  Precipitation*  —  The  precipitation  of  calcium  oxalate 
CaC204,  just  referred  to,  is  a  typical  one  and  may  be  used  to  illus- 
trate the  application  of  ion-product  constancy  (p.  584)  to  explaining 
the  phenomenon.  The  same  explanation  serves  for  all  precipita- 
tions. 

The  first  thing  to  be  remembered  is  that  the  precipitate  which  we 
observe,  however  insoluble  its  material  may.be,  does  not  include  all  of 
the  substance,  but  only  the  excess  beyond  what  is  required  to  saturate 
the  water.  The  liquid  surrounding  the  precipitate  is  always  a  saturated 
solution  of  the  substance  precipitated.  If  it  were  not  so,  some  of  the 
precipitate  would  dissolve  until  the  liquid  became  saturated.  Thus, 
for  example,  when  we  add  ammonium  oxalate  solution  to  calcium 
chloride  solution : 

'  4  (solid), 

the  liquid  is  a  saturated  solution  of  calcium  oxalate,  with  the  excess 
of  this  salt  suspended  in  it. 

Looking  at  the  matter  from  this  view-point,  we  perceive  the  appli- 
cation of  the  rule  of  ion-product  constancy.  In  this  saturated  solu- 
tion (p.  584)  the  product  of  the  ion-concentrations,  [Ca' "]  X  [C204"], 
is  constant.  If  the  original  solutions  had  been  so  very  dilute  that, 
when  they  were  mixed,  the  product  of  the  concentrations  of  these  two 
ions  had  not  reached  the  value  of  this  constant,  no  precipitation  would 
have  occurred.  As  a  matter  of  fact  the  ion-product  considerably  ex- 
ceeded the  requisite  value,  and  hence  the  salt  was  thrown  down  until 
the  balance  remaining  gave  the  value  in  question.  The  rule  for  precip- 
itation, then,  is  as  follows  :  Whenever  the  product  of  the  concentra- 
tions of  any  two  ions  in  a  mixture  exceeds  the  value  of  the  ion-product 
in  a  saturated  solution  of  the  compound  formed  by  their  union,  this 
compound  will  be  precipitated.  Naturally  the  substances  with  small 
solubilities,  and  therefore  small  ion-product  constants,  are  the  ones 
most  frequently  formed  as  precipitates. 

Rule  for  Solution  of  Substances.  —  The  rule  for  solution  of  any 

ionogen  follows  at  once  from  the  foregoing  considerations,  and  may  be 
formulated  by  changing  a  few  of  the  words  in  the  rule  just  given: 
Whenever  the  product  of  the  concentrations  of  any  two  ions  in  a  mix- 
ture is  less  than  the  value  of  the  ion-product  in  a  saturated  solution 
of  the  compound  formed  by  their  union,  this  compound,  if  present  in 


598  INORGANIC   CHEMISTRY 

the  solid  form,  will  be  dissolved.  When  applied  to  the  simplest  case, 
this  rule  means  that  a  substance  will  dissolve  in  a  liquid  not  yet  satu- 
rated with  it,  but  will  not  dissolve  in  a  liquid  already  saturated  with 
the  same  material.  The  value  of  the  rule  lies  in  its  application  to  the 
less  simple,  but  equally  common  cases,  such  as  when  an  insolu- 
ble body  is  dissolved  by  interaction  with  another  substance  (next 
section). 

Applications  of  the  Rule  for  Solution  to  the  Solution  of 
Insoluble  Substances.  —  So  long  as  a  substance  remains  in  pure 
water  its  solubility  is  fixed.  Thus,  with  calcium  hydroxide,  the  sys- 
tem comes  to  rest  when  .17  g.  per  100  C.G.  of  water  have  gone  into 
solution : 

Ca(OH)2  (solid)  <=»Ca  (OH)2  (diss'd)^Ca"  +  20H'. 

But  if  an  additional  reagent  which  can  combine  with  either  one  of  the 
ions  is  added,  the  concentration  of  this  ion  at  once  becomes  less,  the 
ion-product  therefore  tends  to  diminish,  and  further  solution  must  take 
place  to  restore  its  value.  Thus,  if  a  little  of  an  acid  (giving  H')  be 
added  to  the  solution  of  calcium  hydroxide,  the  union  of  OH*  and  H* 
to  form  water  removes  the  OH',  and  solution  of  the  hydroxide  pro- 
ceeds until  the  acid  is  used  up.  There  are  now  more  Ca"  than  OH7 
ions  present,  but  the  ion-product  reaches  the  same  value  as  before,  and 
then  the  change  ceases.  If  a  further  supply  of  acid  is  added,  the 
removal  of  OH'  to  form  H20  begins  again.  With  excess  of  the  acid, 
the  only  stable  OH'  concentration  is  that  which  is  a  factor  in  the  very 
minute  ion-product  of  water,  [OH']  X  [H'].  Hence,  the  calcium 
hydroxide,  which  requires  in  general  a  much  higher  concentration  of 
OH'  than  this  to  precipitate  it  or  to  keep  it  out  of  solution,  finally  all 
dissolves. 

This  particular  action  is  a  neutralization  of  an  insoluble  base. 
But  the  other  kinds  of  actions  by  which  insoluble  ionogens  pass  into 
solution  all  resemble  it  closely,  and  differ  only  in  details.  The  gen- 
eral outlines  of  the  explanation  are  the  same  in  every  case.  We  pro- 
ceed now  to  apply  it  to  the  common  phenomenon  of  the  solution  of  an 
insoluble  salt  by  an  acid. 

Interaction  of  Insoluble  Salts  with  Acids,  Resulting  in 
Solution  of  the  Salt.  —  Calcium  oxalate  passes  into  solution  when 
in  contact  with  acids,  especially  active  acids.  Thus,  with  hydrochloric 


THE  METALS  OF  THE  ALKALINE  EARTHS        599 

acid,  it  gives  calcium  chloride  and  oxalic  acid,  both  of  which  are 
soluble : 

CaC204  f  +  2HC1  <=>  CaCl2  +  H2C204.  (1) 

The  action  of  acids  upon  insoluble  salts  is  so  frequently  mentioned  in 
chemistry  and  is  so  important  a  factor  in  analytical  operations  that  it 
demands  separate  discussion.  This  example  is  a  typical  one  and  may 
be  used  as  an  illustration. 

According  to  the  rules  already  explained  (p.  597),  calcium  oxalate 
(or  any  other  salt)  is  precipitated  when  the  product  of  the  concentra- 
tions of  the  two  requisite  ions  [Ca*']  X  [C204X/]  exceeds  the  ion-pro- 
duct for  a  saturated  solution  of  calcium  oxalate  in  pure  water.  When, 
on  the  contrary,  the  product  of  the  concentrations  of  the  two  ions  falls 
below  the  limiting  value,  a  condition  which  may  arise  from  the 
removal  in  some  way  either  of  the  Ca"  or  of  the  C204"  ions,  the 
undissociated  molecules  will  ionize,  and  the  solid  will  dissolve  to 
replace  them  until  the  ionic  concentrations  necessary  for  equilibrium 
with  the  molecules  have  been  restored  or  until  the  whole  of  the  solid 
present  is  consumed.  Here  the  oxalanion  from  the  calcium  oxalate 
combines  with  the  hydrion  of  the  acid  (usually  an  active  one)  which 
has  been  added,  and  forms  molecular  oxalic  acid : 

C20("  +  2H-  t=t  H2C204.  (2) 

Hence,  dissociation  of  the  dissolved  molecules  of  calcium  oxalate  pro- 
ceeds, being  no  longer  balanced  by  encounters  and  unions  of  the  now 
depleted  ions,  and  this  dissociation  in  turn  leads  to  solution  of  other 
molecules  from  the  precipitate. 

It  will  be  seen  that  the  removal  of  the  ions  in  this  fashion  can 
result  in  considerable  solution  of  the  salt  only  when  the  acid  produced 
is  a  feebly  ionized  one.  Here,  to  be  specific,  the  concentration  of  the 
C2O4"  in  the  oxalic  acid  equilibrium  (2)  above  must  be  less  than  that  of 
the  same  ion  in  a  saturated  calcium  oxalate  solution.  Now  oxalic  acid 
does  not  belong  to  the  least  active  class  of  acids,  and  its  pure  solution 
contains  a  considerable  concentration  of  C2O4".  There  is,  however,  a 
decisive  factor  in  the  situation  which  we  have  not  yet  taken  into 
account.  The  hydrochloric  acid  which  we  used  for  dissolving  the 
precipitate  is  a  very  highly  ionized  acid  and  gives  an  enormously 
greater  concentration  of  hydrion  than  does  oxalic  acid.  Hence  the 
hydrion  is  in  excess  in  equation  (2),  and  the  condition  of  equilibrium 


600  INORGANIC   CHEMISTRY 

[H*i2  v  rc  o  "i 

for    oxalic   acid,  =  —  rTT  p  oN  *       =  ^     w*^  ^e  sa*isned  by  a  corre- 

Li±2L/2u4j 

spondingly  small  concentration  of  C204".  In  this  particular  case, 
therefore,  the  [C2O/]  of  the  oxalic  acid  is  less  than  that  given  by  the 
calcium  oxalate.  The  whole  change,  therefore,  depends  for  its  accom- 
plishment not  only  on  the  mere  presence  of  hydrion,  but  on  the  repres- 
sion of  the  ionization  of  the  oxalic  acid  by  the  great  excess  of  hydrion 
furnished  by  the  active  acid  that  has  been  used.  As  a  matter  of  fact, 
we  find  that  a  weak  acid  like  acetic  acid  has  scarcely  any  effect  upon 
a  precipitate  of  calcium  oxalate.  An  acid  stronger  than  oxalic  acid 
must  be  employed.  The  whole  scheme  of  the  equilibria  is  as  follows  : 

CaC204  (solid)  4=5  CaC204(diss'd)  ±=>Ca"+  C2O4"l 

2HC1  ^2C1'+2H'    f^^jv,0^ 


When  excess  of  an  acid  sufficiently  active  to  furnish  a  large  concentra- 
tion of  hydrion  is  employed,  the  last  equilibrium  is  then  driven  for- 
ward and  the  others  follow.  With  addition  of  a  weak  acid,  only  a 
slight  displacement  occurs,  and  the  system  comes  to  rest  again  when 
the  molecular  oxalic  acid  has  reached  a  sufficient  concentration. 

A  generalization  may  be  based  on  these  considerations  :  an  insol- 
uble salt  of  a  given  acid  will  in  general  interact  and  dissolve  when 
treated  with  a  solution  containing  another  acid,  provided  that  the 
latter  acid  is  a  much  more  highly  ionized  (more  active)  one  than  the 
former  (see  below). 

But  even  active  acids  frequently  fail  to  bring  salts  of  weak  acids 
into  solution,  especially  when  the  weak  acid  is  itself  present  also. 
Here  the  cause  lies  in  the  fact  that  such  salts  are  less  soluble  than 
those  of  the  calcium  oxalate  type,  and  give  so  low  a  concentration  of 
the  negative  ion  that  the  utmost  repression  of  the  ionization  of  the 
corresponding  acid  does  not  give  a  lower  value  for  the  concentration 
of  this  ion  than  does  the  salt  itself.  Thus,  we  have  seen  (p.  375)  that 
even  hydrochloric  acid  (dilute)  will  not  dissolve  a  number  of  sul- 
phides. For  example,  in  the  case  of  cupric  sulphide  in  a  solution 
saturated  with  hydrogen  sulphide,  the  S"  factor  in  the  solubility 
product  [Cu"]  x  [S"]  remains  smaller  than  that  in  the  scheme  defin- 
ing the  hydrogen  sulphide  equilibrium  [H']2  x  [S"]  even  when  the  S" 
factor  in  the  latter  is  diminished  in  consequence  of  great  addition  of 
hydrion.  In  this  case  the  first  link  in  the  chain  of  equilibria: 

CuS  (solid)  ±3  CuS  (diss'd)  *=>  Cu+  S"     1       „  „  M-.,,Hv 
2HC1  £*  201'  +  2H'J  *=*  H2b  (dlss  d)> 


THE    METALS   OF   THE   ALKALINE    EARTHS  601 

tends  so  decidedly  backward  that  only  the  use  of  concentrated  acid  will 
increase  the  concentration  of  the  H"  to  an  extent  sufficient  to  secure 
any  marked  advance  of  the  whole  action.  We  must  add,  there- 
fore, to  the  above  rule:  provided  also  that  the  salt  is  not  one  of 
extreme  insolubility.  This  point  will  be  illustrated  more  fully  in 
connection  with  the  description  of  individual  sulphides  (see  under 
Cadmium). 

Illustrations  of  the  application  of  these  generalizations  are  count- 
less. Carbonic  acid  is  made  from  marble  (p.  480),  hydrogen  sulphide 
from  ferrous  sulphide  (p.  371),  hydrogen  peroxide  from  sodium  per- 
oxide (p.  303),  and  phosphoric  acid  from  calcium  phosphate  (p.  456). 
In  each  case  the  acid  employed  to  decompose  the  salt  is  more  active 
than  the  acid  to  be  liberated.  On  the  other  hand,  calcium  oxalate  is 
insoluble  in  acetic  acid  because  this  acid  is  weaker  than  is  oxalic  acid. 
We  have  thus  only  to  examine  the  list  of  acids  showing  their  degrees 
of  ionization  (p.  330)  in  order  to  be  able  to  tell  which  salts,  if  insol- 
uble in  water,  will  be  dissolved  by  acids,  and,  in  general,  what  acids 
will  be  sufficiently  active  in  each  case  for  the  purpose.  In  chemical 
analysis  we  discriminate  between  salts  soluble  in  water,  those  soluble 
in  acetic  acid  (the  insoluble  carbonates  and  some  sulphides,  FeS  and 
ZnS,  for  example),  those  requiring  active  mineral  acids  for  their  solu- 
tion (calcium  oxalate  and  the  more  insoluble  sulphides,  for  example), 
and  those  insoluble  in  all  acids  (barium  sulphate  and  other  insoluble 
salts  of  active  acids). 

Precipitation   of  Insoluble   Salts  in   Presence  of  Acids. — 

The  converse  of  solution,  namely,  precipitation,  depends  upon  the 
same  conditions  :  an  insoluble  salt  which  is  dissolved  by  a  given  acid 
cannot  be  formed  by  precipitation  in  the  presence  of  this  acid.  Thus, 
calcium  oxalate  can  be  precipitated  in  presence  of  acetic  acid,  but  not 
in  presence  of  active  mineral  acids  in  ordinary  concentrations.  Cupric 
sulphide  or  barium  sulphate  can  be  precipitated  in  presence  of  any  acid, 
but  ferrous  sulphide  and  calcium  carbonate  only  in  the  absence  of 
acids. 

From  this  it  does  not  follow  that  calcium  oxalate,  for  example, 
cannot  be  precipitated  if  once  an  active  acid  has  been  added  to  the 
mixture.  To  secure  precipitation,  all  that  is  necessary  is  to  remove 
the  excess  of  hydrion  which  is  repressing  the  ionization  of  the  oxalic 
acid.  This  can  be  done  by  adding  a  base,  which  removes  the  H*,  or 
even  by  adding  sodium  acetate.  The  acetanion  CjJLjO/  unites  with 


602  INORGANIC   CHEMISTRY     , 

the  H*  to  form  the  little  ionized  acetic  acid,  in  presence  of  which 
calcium  oxalate  can  be  precipitated. 

Calcium  Carbide.  —  The  manufacture  of  this  compound  has  been 
described  (p.  478),  and  the  formation  of  acetylene  by  its  interaction 
with  water  has  already  been  discussed  (p.  496).  The  substance  was 
discovered  by  Wohler  in  1862,  was  first  prepared  by  the  use  of  elec- 
trical heating  by  Borchers  in  1891,  and  was  made  on  a  large  scale  in 
1892  by  Wilson,  a  Canadian  engineer. 

Bleaching  Powder.  —  This  substance  (cf.  p.  266)  is  manufactured 
by  conducting  chlorine  into  a  box-like  structure  containing  slaked  lime 
spread  upon  perforated  shelves.  When  the  transformation  is  complete, 
the  supply  of  the  gas  is  shut  off,  and  some  lime-dust  is  blown  into  the 
chamber  to  absorb  the  remainder  of  the  free  chlorine.  The  action  is 
represented  by  the  equation  already  given,  or  by  the  following : 

2Ca(OH)2  +  2CL,  ->  CaCl2  +  Ca(OCl)2  +  2H2O. 

While  pure  lime  should  thus  yield  a  product  containing  49  per  cent  of 
chlorine,  in  practice  the  proportion  is  always  less.  Good  bleaching 
powder  should  contain  36-37  per  cent  of  chlorine. 

That  bleaching  powder  is  a  mixed  salt  CaCl(ClO)  rather  than  an 
equi-molar  mixture  of  calcium  chloride  and  calcium  hypochlorite, 
which  would  have  the  same  composition,  is  rendered  probable  by  the 
facts  that  the  material  is  not  deliquescent  as  is  calcium  chloride,  and 
that  calcium  chloride  cannot  be  dissolved  out  of  it  by  alcohol. 

Bleaching  powder  is  somewhat  soluble  in  water,  and  in  solution 
the  ions  Ca",  Cl',  and  CIO'  are  all  present.  Addition  of  acids  causes 
the  formation  of  hydrochloric  and  hypochlorous  acids.  The  oxidizing 
and,  incidentally,  the  bleaching  properties  (p.  269)  of  the  latter  are 
characteristic  of  the  acidified  liquid.  Weak  acids  like  carbonic  acid 
displace  the  hypochlorous  acid  only  (cf.  p.  267),  and  hence  the  dry 
powder,  when  exposed  to  the  air,  has  the  odor  of  the  latter  substance 
rather  than  that  of  chlorine. 

The  substance  is  largely  used  by  bleachers  (cf.  p.  270),  and  as  a 
disinfectant  to  destroy  germs  of  putrefaction  and  disease. 

Calcium  Nitrate.  —  This  salt  is  found  in  the  soil  (p.  438),  and 
may  best  be  prepared  in  pure  form  by  treating  marble  with  nitric  acid 
and  allowing  the  ^product  to  crystallize  from  the  solution.  Calcium 


THE  METALS  OF  THE  ALKALINE  EARTHS        603 

nitrate  forms  several  hydrates.  The  tetrahydrate  Ca(N08)2,4H2OJ 
which  forms  transparent  monoclinic  prisms,  is  the  one  deposited  at 
ordinary  temperatures.  The  anhydrous  salt  is  easily  soluble  in  alco- 
hol. It  is  used  in  the  laboratory  for  drying  nitrogen  peroxide.  When 
heated  it  decomposes  (cf.  p.  444),  giving  nitrogen  peroxide,  oxygen, 
and  quicklime. 

Calcium  Sulphate.  —  This  salt  is  found  in  large  quantities  in 
nature.  The  mineral  anhydrite  CaS04  occurs  in  the  salt  layers  (see 
under  Manganous  sulphate).  It  contains  no  water  of  crystallization, 
and  its  crystals  belong  to  the  rhombic  system.  The  dihydrate, 
CaS04,2H20,  is  more  plentiful.  In  granular  masses  it  constitutes  ala- 
baster. When  perfectly  crystallized  it  is  named  gypsum  or  selenite. 
The  same  hydrate  is  formed  by  precipitation  from  solutions.  Its 
solubility  is  about  1  in  500  at  18°.  Its  solubility  varies  in  an  unusual 
manner  with  temperature,  increasing  slowly  to  38°  and  then  falling 
off. 

When  its  temperature  is  raised,  the  dihydrate  quickly  shows  an 
appreciable  aqueous  tension.  After  three-fourths  of  the  water  has 
escaped,  a  definite  hydrate  (CaS04)2,  H20  remains.  This  hemihydrate 
shows  a  much  smaller  tension  of  water  vapor  (cf.  p.  122). 

The  transition  temperature  at  which  the  dihydrate  passes  sharply  into  the  hemi- 
hydrate is  107°.  It  corresponds  to  the  temperature  of  35.2°  at  which  the  decahy- 
drate  of  sodium  carbonate  turns  into  the  monohydrate  and  water.  At  107°  both 
of  the  hydrates  are  in  equilibrium  with  water  (p.  594).  Naturally  this  state  of 
affairs  can  be  realized  only  in  a  tube  sealed  up  to  prevent  the  escape  of  the  water. 

Plaster  of  paris  is  manufactured  by  heating  gypsum  until  nearly 
all  the  water  of  hydration  has  been  driven  out.  When  it  is  mixed 
with  water,  the  dihydrate  is  quickly  re-formed  and  a  rigid  mass  is  pro- 
duced. If,  in  course  of  manufacture,  the  water  is  all  removed,  or  the 
temperature  is  allowed  to  rise  much  above  the  most  favorable  one 
(about  125°),  the  product  when  mixed  with  water  does  not  set  quickly 
and  is  said  to  be  "  dead-burnt."  In  explanation  of  this  it  should  be 
noted  that  natural  anhydrite  combines  very  slowly  with  water. 
Apparently  good  plaster  of  paris  must  contain  some  unchanged  parti- 
cles of  the  dihydrate  which  may  act  as  nuclei.  They  fulfil  the  same 
r61e  as  the  crystal  which  is  added  to  a  supersaturated  solution  (p.  159), 
without  which  crystallization  may  be  long  delayed  or  may  even  fail  to 
take  place.  Probably  with  moderate  heating  the  product  is  a  mixture 
of  the  dihydrate  and  the  hemihydrate  with  anhydrous  salt,  while  the 


604  INORGANIC   CHEMISTRY 

more  rapid  decomposition  at  higher  temperatures  destroys  all  of  the 
first.  The  former  mixture  must  be  an  unstable  system,  and  the  dihy- 
drate  loses  water  to  the  anhydrous  salt.  At  ordinary  temperatures, 
however,  this  transference  must  be  very  slow,  and  hence  the  property 
of  setting  is  not  lost  by  prolonged  storage. 

That  the  plaster  sets,  instead  of  forming  a  loose  mass  of  dihydrate, 
is  due  to  the  fact  that  the  anhydrous  salt  is  more  soluble  than  the 
dihydrate,  and  so  a  constant  solution  of  the  one  and  deposition  of  the 
other  goes  on  until  the  hydration  is  complete  : 

CaSO,  (solid)  ^  CaSO  (diss'd) )  ^  ^^  (soM). 

2  / 

This  process  results  in  the  formation  of  an  interlaced  and  coherent 
mass  of  minute  crystals. 

Plaster  of  paris  is  used  for  making  casts  and  in  surgery.  The 
setting  of  the  material  is  accompanied  by  a  slight  increase  in  volume, 
and  hence  a  very  sharp  reproduction  of  all  the  details  in  the  structure 
of  the  mold  is  obtained.  An  "  ivory  "  surface,  which  makes  washing 
practicable,  is  conferred  by  painting  the  cast  with  a  solution  of  paraffin 
or  stearine  in  petroleum  ether.  The  waxy  material,  left  by  evaporation 
of  the  volatile  hydrocarbons,  fills  the  pores  and  prevents  solution  and 
disintegration  of  the  substance  by  water.  Stucco  is  made  with  plaster 
of  paris  and  rubble,  and  is  mixed  with  a  solution  of  size  or  glue  in- 
stead of  water. 

Calcium,  Sulphide.  —  This  compound  is  most  easily  made  by 
strongly  heating  pulverized  calcium  sulphate  and  charcoal.  The 
sulphate  is  reduced  : 

4C  +  CaS04  -»  CaS  +  4CO. 

Calcium  sulphide  is  meagerly  soluble  in  water,  but  is  nevertheless 
slowly  dissolved  in  consequence  of  its  decomposition  by  hydrolysis 
into  calcium  hydroxide  and  calcium  hydrosulphide  (cf.  p.  375).  It  is 
probable  that  the  action  would  be  less  nearly  complete  than  it  is  if 
the  reverse  action  were  not  weakened  by  the  precipitation  of  the  cal- 
cium hydroxide  : 

2CaS  +  2H20  <=>  Ca(OH)2  J+  Ca(SH)2. 

Since  calcium  sulphide  is  thus  decomposed  by  water  it  cannot  be 
precipitated  from  aqueous  solution  by  adding  a  soluble  sulphide,  such 


THE  METALS  OF  THE  ALKALINE  EARTHS        605 

as  ammonium  sulphide,  to  a  solution  of  a  salt  of  calcium.     Only  the 
soluble  hydrosulphide  can  be  formed. 

Ordinary  calcium  sulphide,  after  it  has  been  exposed  to  sunlight, 
usually  shines  in  the  dark.  Barium  sulphide  behaves  in  the  same  way. 
On  this  account  these  substances  are  used  in  making  luminous  paint. 
They  apparently  owe  this  behavior  to  the  presence  of  traces  of  com- 
pounds of  vanadium  and  bismuth,  for  the  purified  substances  are  not 
affected  in  the  same  fashion. 

Phosphates  of  Calcium.  —  The  tertiary  orthophosphate  of  cal- 
cium Ca3(P04)2,  known  as  phosphorite,  is  found  in  many  localities.  It 
is  probably  derived  from  the  remains  of  animals.  Guano  contains 
some  of  the  same  substance,  along  with  compounds  of  nitrogen. 
Apatite,  3Ca3(P04)2,CaF2,  a  double  salt  with  calcium  fluoride,  is  a 
common  mineral  and  frequent  component  of  rocks.  The  orthophos- 
phate forms  about  83  per  cent  of  bone-ash,  and  is  contained  also  in 
the  ashes  of  plants.  It  may  be  precipitated  by  adding  a  soluble 
phosphate  to  a  solution  of  a  salt  of  calcium. 

Since  it  is  a  salt  of  a  weak  acid,  and  belongs  to  the  less  insoluble 
class  of  such  salts,  calcium  phosphate  is  dissolved  by  dilute  mineral 
acids  (cf.  p.  598),  the  ions  HP04"  and  H2PO/  being  formed.  When 
a  base,  such  as  ammonium  hydroxide,  is  added  to  the  solution,  the 
calcium  phosphate  is  reprecipitated  (cf.  p.  601). 

Calcium  phosphate  is  chiefly  used  in  the  manufacture  of  phos- 
phorus and  phosphoric  acid  (p.  456),  and  as  a  fertilizer.  The  supply 
of  calcium  phosphate  in  the  soil  arises  from  the  decomposition  of  rocks 
containing  phosphates,  and  is  gradually  exhausted  by  the  removal 
of  crops.  Bone-ash  is  sometimes  used  to  make  up  the  deficiency. 
It  is  almost  insoluble  in  water,  however,  and,  although  somewhat 
less  insoluble  in  natural  water  containing  salts  like  sodium  chloride 
(cf.  p.  601),  is  brought  into  a  condition  for  absorption  by  the  plants 
too  slowly  to  be  of  much  service.  In  consequence  of  this  the  "  super- 
phosphate "  (see  below)  is  preferred. 

Primary  calcium  orthophosphate  is  manufactured  in  large  quan- 
tities from  phosphorite  by  the  action  of  sulphuric  acid.  The  uncon- 
centrated  "  chamber  acid "  is  used  for  this  purpose,  as  water  is  re- 
quired in  the  resulting  action.  The  amounts  of  material  employed 
correspond  to  the  equation  : 

Ca,(P04)2  +  2H2S04  +  6HaO  ->  Ca(HaPO4)2,2H2O  +  2CaSO4,2H2O. 


606  INORGANIC   CHEMISTRY 

As  soon  as  mixture  has  been  effected,  the  action  proceeds  with  evolu- 
tion of  heat,  and  a  large  cake  of  the  two  hydrated  salts  remains.  This 
mixture,  after  being  broken  up,  dried,  and  packed  in  bags,  is  sold  as 
«  superphosphate  of  lime."  The  primary  phosphate  which  it  contains 
is  readily  soluble  in  water,  and  is  therefore  of  great  value  as  a  fer- 
tilizer. 

Calcium  Silicate.  —  Calcium  metasilicate  CaSiO3  forms  the  min- 
eral wollastonite,  which  is  rather  scarce,  and  enters  into  the  composi- 
tion of  many  complex  minerals,  such  as  garnet,  mica,  and  the  zeolites. 
It  may  be  made  by  precipitation  with  a  solution  of  sodium  metasilicate 
(p.  577),  or  by  fusing  together  powdered  quartz  and  calcium  carbonate 
or  quicklime  : 

Si02  +  CaC03  ->  CaSiO8  +  C02. 

Glass.  —  Common  glass  is  a  complex  silicate  of  sodium  and  cal- 
cium, or  a  homogeneous  mixture  of  the  silicates  of  these  metals  with 
silica.  It  has  a  composition  represented  approximately  by  the  formula 
Na2O,CaO,  6Si02,  and  is  made  by  melting  together  sodium  carbonate, 
limestone,  and  pure  sand  : 


Na2C08  +  CaCO8  +  6Si02  ->  Na^OjCaOjGSiO,  +  2CO2. 

For  the  most  fusible  glass,  a  smaller  proportion  of  quartz  is  employed. 
This  variety  is  known  as  soda-glass,  or,  from  its  easy  fusibility,  as 
soft  glass.  First,  the  materials  are  heated  to  a  temperature  high 
enough  to  produce  chemical  action  without  bringing  about  complete 
melting.  This  permits  the  ready  escape  of  the  gases.  Then  the  tem- 
perature is  raised  to  about  1200°  until  fusion  is  complete  and  all  the 
bubbles  have  escaped.  Finally,  the  crucible  and  its  contents  are 
allowed  to  cool  to  700-800°  to  allow  the  latter  to  acquire  the  viscosity 
required  for  working. 

Plate-glass  is  made  by  casting  the  material  in  large  sheets  and 
polishing  the  surfaces  until  they  are  plane.  Window-glass  is  prepared 
by  blowing  bulbs  of  long  cylindrical  shape,  and  ripping  them  down  one 
side  with  the  help  of  a  diamond.  The  resulting  curved  sheets  are  then 
placed  on  a  flat  surface  in  a  furnace  and  are  there  allowed  to  open  out. 
Beads  are  made,  chiefly  in  Venice,  by  cutting  narrow  tubes  into  very  short 
sections  and  rounding  the  sharp  edges  by  fire.  Ordinary  apparatus  is 
made  of  soft  soda-glass,  and  hence  when  heated  strongly  it  tends  to 
soften  and  also  to  confer  a  strong  yellow  tint  (cf.  p.  563)  on  the  flame. 


THE  METALS  OF  THE  ALKALINE  EARTHS       607 

In  all  cases  the  articles  are  annealed  by  being  passed  slowly  through 
a  special  furnace  in  which  their  temperature  is  lowered  very  grad- 
ually. Glass  which  has  been  suddenly  chilled  is  in  a  state  of  tension 
and  breaks  easily  when  handled. 

Bottles  are  made  with  impure  materials,  and  owe  their  color  chiefly 
to  the  silicate  of  iron  which  they  contain.  In  making  cheap  glass, 
sodium  sulphate  is  often  substituted  for  the  much  more  expensive 
carbonate.  In  this  case  powdered  charcoal  or  coal  is  added  to  reduce 
the  sulphate  : 

2Na2S04  +  C  +  2Si02  ->  2Na^Si08  +  C02  +  2S02. 


Soft  glass  is  partially  dissolved  by  water.  When  powdered  glass 
is  shaken  with  water  the  solution  soon  dissolves  enough  sodium  silicate 
to  give  the  alkaline  reaction  (pink  color)  with  phenolphthalein  (of. 
p.  355). 

Bohemian,  or  hard  glass,  is  much  more  difficult  to  fuse  than  soda- 
glass,  and  is  also  much  less  soluble  in  water.  It  is  manufactured  by 
substituting  potassium  carbonate  for  the  sodium  carbonate,  and  is  used 
for  making  apparatus  for  special  purposes  where  infusibility  or  insolu- 
bility are  desirable.  When  lead  oxide  is  employed  instead  of  lime- 
stone, a  soda-lead  glass  known  as  flint  glass  is  produced.  This  has  a 
high  specific  gravity,  and  a  great  refracting  power  for  light,  and  is 
employed  for  making  glass  ornaments.  By  the  use  of  grinding 
machinery,  cut  glass  is  made  from  it.  It  is  easily  fusible,  and  dissolves 
in  water  like  soda-glass. 

Colored  glass  is  prepared  by  adding  small  amounts  of  various 
oxides  to  the  usual  materials.  The  oxides  combine  with  the  silica,  and 
produce  strongly  colored  silicates.  Thus,  cobalt  oxide  gives  a  blue, 
chromium  oxide  or  cupric  oxide  a  green,  and  uranium  a  yellow  glass. 
Cuprous  oxide,  with  a  reducing  agent,  and  compounds  of  gold,  give  the 
free  metals,  suspended  in  colloidal  solution  in  the  glass,  and  confer  a 
deep-red  color  upon  it.  Milk-glass  contains  finely  powdered  calcium 
phosphate  in  suspension,  and  white  enamels  are  made  by  adding 
stannic  oxide. 

Glass  is  a  typical  amorphous  substance  (cf.  p.  139).  From  the  facts 
that  it  has  no  crystalline  structure,  and  that  it  softens  gradually  when 
warmed,  instead  of  showing  a  definite  melting-point,  it  is  regarded  as 
a  supercooled  liquid  of  extreme  viscosity.  Most  single  silicates  crys- 
tallize easily,  and  have  definite  freezing-  (and  melting-)  points.  Glass 
may  be  regarded  as  a  solution  of  several  silicates.  When  kept  for  a 


608  INORGANIC   CHEMISTRY 

considerable  length  of  time  at  a  temperature  insufficient  to  render  it 
perfectly  fluid,  some  of  its  components  crystallize  out,  the  glass  becomes 
opaque,  and  "  devitrification  "  is  said  to  have  occurred.  The  absence  of 
such  changes  in  cold  glass  may  be  attributed  to  that  general  hampering 
of  all  molecular  movements  and  interactions  which  is  characteristic  of 
low  temperatures.  The  word  "  crystal "  popularly  applied  to  glass  is 
thus  definitely  misleading. 

Calcion :  Analytical  Reactions.  —  Ionic  calcium  is  colorless.  It 
is  bivalent,  and  combines  with  negative  ions.  Many  of  the  resulting 
salts  are  more  or  less  insoluble  in  water.  Upon  the  insolubility  of 
the  carbonate,  phosphate,  and  oxalate  are  based  tests  for  calcion  in 
qualitative  analysis  (see  below).  The  presence  of  the  element  is  most 
easily  recognized  by  the  brick-red  color  its  compounds  confer  on  the 
Bunsen  flame,  and  by  two  bands  —  a  red  and  a  green  one  —  which  are 
shown  by  the  spectroscope  (p.  561). 

STRONTIUM. 

The  compounds  of  strontium  resemble  closely  those  of  calcium, 
both  in  physical  properties  and  in  chemical  behavior. 

Occurrence'  —  The  carbonate  of  strontium  SrC08  is  found  as 
strontianite  (Strontian,  a  village  in  Argyleshire),  and  is  isomorphous 
with  aragonite.  The  sulphate,  celestite  SrSO4,  is  more  plentiful.  It 
shows  rhombic  crystals,  which  are  isomorphous  with  those  of  anhydrite, 
often  have  a  blue  color,  and  are  commonly  associated  with  native  sul- 
phur in  specimens  from  Sicily.  The  metal  may  be  isolated  by  elec- 
trolysis of  the  molten  chloride. 

Compounds  of  Strontium.  —  The  compounds  are  all  made  from 
the  natural  carbonate  or  sulphate.  The  former  may  be  dissolved 
directly  in  acids,  and  the  latter  is  first  reduced  by  means  of  carbon  to 
the  sulphide,  and  then  treated  with  acids. 

Strontium  chloride,  made  in  one  of  the  above  ways,  is  deposited 
from  solution  as  the  hexahydrate.  The  nitrate  comes  out  of  hot  solu^ 
tions  in  octahedrons  which  are  anhydrous.  From  cold  water  the  tet- 
rahydrate  is  obtained  (see  under  Manganous  sulphate).  The  anhy- 
drous salt  is  mixed  with  sulphur,  charcoal,  and  potassium  chlorate  to 
make  «  red  fire."  The  oxide  SrO  may  be  secured  by  igniting  the  car- 


THE  METALS  OF  THE  ALKALINE  EARTHS        609 

bonate,  but  on  account  of  the  low  dissociation  tension  of  the  compound 
it  is  obtained  with  greater  difficulty  than  is  calcium  oxide  from  cal- 
cium carbonate.  It  is  generally  made  by  heating  the  nitrate  or 
hydroxide. 

Strontium  hydroxide  is  made  by  heating  the  carbonate  in  a 
current  of  superheated  steam: 

SrCO8  +  H20  ->  Sr(OH)2  +  C02. 

This  action  takes  place  more  easily  than  does  the  mere  dissociation  of 
the  carbonate,  because  the  formation  of  the  hydroxide  liberates  energy, 
and  this  partially  compensates  for  the  energy  which  has  to  be  pro- 
vided to  decompose  the  carbonate  (cf.  p.  78).  The  lowering  of  the 
partial  pressure  of  the  carbon  dioxide  by  the  steam  also  contributes  to 
the  result  (cf.  p.  595). 

The  hydrate  crystallizes  from  water  as  Sr(OH)2, 8H20,  and  is 
employed  in  separating  crystallizable  sugar  from  molasses.  By  evapo- 
ration of  the  extract  from  the  sugar-cane  or  beet-root,  as  much  of  the 
sugar  as  possible  is  first  secured  by  crystallization.  Then  the  molasses 
which  remains  is  mixed  with  a  saturated  solution  of  strontium  hydrox- 
ide. The  resulting  precipitate  of  sucrate  of  strontium,  C^H^On, 
SrO,  or  CjjjHgjjOj^SrO,  is  separated  by  a  filter-press,  made  into  a  paste 
with  water,  and  treated  with  carbon  dioxide.  A  second  filtration  parts 
the  insoluble  carbonate  of  strontium  from  the  solution  of  sugar,  and  the 
latter  is  evaporated  and  allowed  to  crystallize.  Calcium  hydroxide, 
which  gives  a  tricalcium  sucrate,  is  often  employed  in  the  same  way. 

Strontion  is  a  bivalent  ion,  and  gives  insoluble  compounds  with 
carbonanion,  sulphanion,  and  oxalanion.  The  presence  of  strontium  is 
recognized  by  the  carmine-red  color  which  its  compounds  give  to  the 
Bunsen  flame  (see  also  below).  Its  spectrum  shows  several  red  bands 
and  a  very  characteristic  blue  line. 

BARIUM. 

The  physical  and  chemical  properties  of  the  compounds  of  barium 
recall  those  of  strontium  and  calcium.  All  the  compounds  of  barium 
which  are  soluble  in  water,  or  can  be  brought  into  solution  by  the  weak 
acids  of  the  digestive  fluids,  are  poisonous. 

Occurrence.  —  Like  strontium,  barium  is  found  in  the  form  of 
the  carbonate,  witherite  BaCO8,  which  is  rhombic  and  isomorphous  with 


610  INORGANIC   CHEMISTRY 

aragonite,  and  the  sulphate  BaS04,  heavy-spar  or  barite  (Gk. 
heavy),  which  is  rhombic  and  isomorphous  with  anhydrite.  The 
specific  gravity  of  the  sulphate  is  4.5,  while  the  specific  gravity  of  other 
compounds  of  the  light  metals  does  not  generally  exceed  2.5.  The  free 
metal,  which  is  silver-white,  may  be  obtained  by  electrolysis  of  the 
molten  chloride. 

The  compounds  are  made  by  treating  the  natural  carbonate  with 
acids  directly,  or  by  first  reducing  the  sulphate  with  carbon  to  sulphide, 
or  converting  the  carbonate  into  oxide,  and  then  treating  the  products 
with  acids. 

Barium  Carbonate.  —  This  carbonate  demands  so  high  a  tempera- 
ture (about  1500°)  for  the  attainment  of  a  sufficient  dissociation  ten- 
sion, and  is  so  apt  then  to  be  partially  protected  from  decomposition 
by  the  melting  of  the  oxide,  that  special  means  is  employed  for  its 
decomposition.  It  is  heated  with  powdered  charcoal  (cf.  p.  485)  : 

BaC08  +  C  ->  BaO  +  2CO. 

The  precipitated  form  of  the  carbonate  is  made  by  adding  sodium  car- 
bonate to  the  aqueous  extract  from  crude  barium  sulphide  (q.v.).  The 
compound  is  also  obtained  by  fusing  pulverized  barite  with  excess  of 
sodium  carbonate,  and  dissolving  the  sodium  salts  out  of  the  residue. 

The  Sulphate  and  Sulphide.  —  The  natural  sulphate  is  the 
source  of  many  of  the  compounds  of  barium.  The  precipitated  sul- 
phate, made  by  adding  sulphuric  acid  to  the  aqueous  extract  from 
barium  sulphide,  is  used  in  making  white  paint,  in  filling  paper  for 
glazed  cards,  and  sometimes  as  an  adulterant  of  white  lead.  The  salt  is 
highly  insoluble  in  water  and  is  hardly  at  all  affected  by  aqueous  solu- 
tions of  chemical  agents.  It  is  somewhat  soluble  in  hot,  concentrated 
sulphuric  acid,  and  the  solution  yields  crystals  of  a  compound 
BaS04,  H2S04,  or  Ba(HS04)2.  Calcium  and  strontium  sulphates  behave 
in  the  same  way.  All  three  compounds  are  decomposed  by  water,  and 
give  the  insoluble  sulphates. 

Barium  sulphide,  like  the  sulphides  of  calcium  and  strontium 
(p.  604),  is  slightly  soluble  in  water,  but  slowly  passes  into  solution 
owing  to  hydrolysis  and  formation  of  the  hydroxide  and  hydrosulphide. 
It  is  made  by  heating  the  pulverized  sulphate  with  charcoal : 

BaS04  +  40  ->  BaS  +  4CO. 


THE  METALS  OF  THE  ALKALINE  EARTHS        611 

The  Chloride  and  Chlorate.  —  Barium  chloride  is  generally 
manufactured  by  heating  the  sulphide  with  calcium  chloride.  The 
whole  treatment  of  the  heavy-spar  is  carried  out  in  one  operation  : 

BaS04  +  4C  +  CaCl2  -*  4CO  +  BaCL,  +  CaS. 

By  rapid  treatment  with  water,  the  chloride  can  be  separated  from  the 
calcium  sulphide  before  much  decomposition  of  the  latter  (cf.  p.  604) 
has  taken  place.  Barium  chloride  crystallizes  in  rhombic  tables  as  a 
dihydrate  BaCl2, 2H20.  Aside  from  the  difference  in  composition,  this 
compound  differs  from  the  ordinary  hydrated  chlorides  of  calcium  and 
strontium  in  being  non-hygroscopic  and  in  being  capable  of  dehydration 
by  heat  without  the  formation  of  any  hydrogen  chloride  (cf.  p.  535). 

Barium  chlorate  is  made  by  treating  the  precipitated  barium  car- 
bonate with  a  solution  of  chloric  acid.  It  is  deposited  in  beautiful 
monoclinic  crystals,  and  is  used  with  sulphur  and  charcoal  in  the  prepa- 
ration of  "  green  fire." 

The  Oxides  and  Hydroxide.  —  Oxide  of  barium  BaO  is  manufac- 
tured from  the  carbonate  or  sulphide.  In  the  latter  case,  moist  carbon 
dioxide  is  passed  over  the  sulphide,  and  the  resulting  carbonate  is  then 
treated  with  steam.  The  compound  may  be  obtained  in  pure  form  by 
heating  the  nitrate.  The  oxide  unites  with  water  to  form  the  hydrox- 
ide Ba(OH)2.  When  heated  in  a  stream  of  air  or  oxygen  it  gives  the 
dioxide  Ba02.  This  change  and  its  reversal  constitute  the  basis  of 
Brin's  process  for  obtaining  oxygen  from  the  air  (p.  63).  To  protect 
the  oxide  from  conversion  into  the  carbonate  and  hydrate,  which  are 
not  decomposable  at  the  temperature  employed,  the  air  must  be  care- 
fully purified  from  carbon  dioxide  and  moisture. 

Barium  dioxide,  when  made  by  union  of  oxygen  with  the  monoxide, 
is  a  compact  gray  mass.  A  hydrated  form  is  thrown  down  as  a  crystal- 
line precipitate  when  hydrogen  peroxide  solution  is  added  to  a  solution 
of  barium  hydroxide  : 

Ba(OH)2  +  H202  ±?  Ba02  J  +  2H2O. 

The  crystals  have  the  formula  Ba02, 8H2O.  Similar  hydrates  of  the 
dioxides  of  strontium  and  calcium  may  be  made  in  the  same  way. 
In  all  three  cases  the  pure  dioxides  are  obtained  as  white  powders  by 
removal  of  the  water  of  hydration  by  very  gentle  heating  in  vacua 
(cf.  p.  276).  The  dioxides  of  strontium  and  calcium  are  not  formed  by 
direct  union  of  oxygen  with  the  oxides.  Barium  dioxide  is  used  in 
the  manufacture  of  hydrogen  peroxide  (p.  303). 


612  INORGANIC   CHEMISTRY 

Barium  hydroxide  is  the  most  soluble  of  the  hydroxides  of  this 
group,  and  gives,  therefore,  the  highest  concentration  of  hydroxidion. 
The  solution  is  known  as  "  baryta- water."  It  is  also  the  most  stable  of 
the  three  hydroxides,  and  may  be  melted  without  decomposition.  It 
crystallizes  when  a  warm,  saturated  solution  cools  in  the  form 
Ba(OH)2, 8H20.  It  is  much  used  in  quantitative  analysis  for  making 
standard  alkali-solutions.  Solutions  of  sodium  or  potassium  hydrox- 
ide may  acquire  varying  proportions  of  carbonate  by  the  action  of 
carbon  dioxide  from  the  air,  and  their  action  on  indicators  loses  there- 
by in  sharpness.  With  barium  hydroxide  this  is  impossible,  for  the 
carbonate  is  insoluble,  and  is  precipitated  from  the  solution. 

Barium  Nitrate.  —  This  salt  is  made  by  the  action  of  nitric  acid 
on  the  sulphide,  oxide,  hydroxide,  or  carbonate  of  barium.  It  crys- 
tallizes from  aqueous  solution  without  water  of  hydration. 

Analytical  Reactions  of  the  Calcium  Family.  —  Barion  is  a 
colorless,  bivalent  ion.  Many  of  its  compounds  are  insoluble  in  water, 
and  the  sulphate  is  insoluble  in  acids  also.  The  spectrum  given  by 
the  salts  contains  a  number  of  green  and  orange  lines. 

In  solutions  of  salts  of  calcium,  strontium,  and  barium,  the  ions  cal- 
cion,  strontion,  and  barion  may  be  distinguished  by  the  fact  that 
calcium  sulphate  solution  will  precipitate  the  strontium  and  barium 
as  sulphates,  but  will  leave  salts  of  calcium  unaffected.  Similarly, 
strontium  sulphate  solution  precipitates  barium  sulphate,  and  does 
not  give  any  result  with  salts  of  the  two  first.  The  oxalate  of  calcium 
is  precipitated  in  presence  of  acetic  acid,  while  the  oxalates  of 
strontium  and  barium  are  not  (cf.  p.  599),  and  there  are  other  differ- 
ences of  a  like  nature  in  the  solubilities  of  the  salts. 

Exercises.  —  1.  Arrange  the  chromates  of  the  metals  of  this 
family  in  the  order  of  solubility  (p.  544).  Compare  the  solubilities 
with  those  of  the  carbonates,  oxalates,  and  sulphates  of  the  metals  of 
the  same  family. 

2.  What  must  be  the  approximate  total  molar  concentration  of  the 
solution  of  calcium  chloride  freezing  at  —  48°  (p.  291)  ? 

3.  What  is  meant  by  fluorescence  (cf.  any  book  on  physics)  ? 

4.  What  will  be  the  ratio  by  volume,  at  150°,  of  the  nitrogen  per- 
oxide and  oxygen  given  off  by  the  decomposition  of  calcium  nitrate  ? 


THE  METALS  OF  THE  ALKALINE  EARTHS        613 

What  would  be  the  nature  of  the  difference  between  the  ratio  at  150° 
and  that  at  room  temperature  ? 

5.  What  fact  about  the  heat  of  solution  of  gypsum  is  indicated  by 
its  change  of  solubility  with  temperature  (p.  260)  ? 

6.  What   is   the   significance   of   the   fact  that  hydrated  barium 
chloride  gives  no  hydrogen  chloride  when  heated  ? 

7.  What  are  the  advantages  of  removing  water  of  hydration  in 
vacua  (p.  611)  ? 

8.  Explain  in  terms  of  the  ionic  hypothesis  the  precipitation  of 
the  sulphate  of  strontium  by  calcium  sulphate  solution,  and  the  ab- 
sence of  precipitation  when  the  latter  is  added  to  the  solution  of  a 
soluble  salt  of  calcium. 

9.  Construct  a  table  for  the  purpose  of  comparing  the  properties  of 
the  free  elements  of  this  family  and  also  the  properties  of  their  corre- 
sponding compounds. 

10.  Are  the  elements  of  this  family  typical  metals?     If  not,  in 
what  respects  do  they  fall  short  (p.  533)  ? 

11.  What  inference  do  you  draw  from  the  fact  that  the  oxalates 
of  barium  and  strontium  are  not  precipitated  in  presence  of   acetic 
acid,  while  the  oxalate  of  calcium  is  so  precipitated  ?     Is  the  infer- 
ence confirmed  by  reference  to  the  data  ? 


CHAPTER    XXXVI 


COPPER,  SILVER,  GOLD 

THE  three  metals  of  this  family,  being  found  free  in  nature,  are 
amongst  those  which  were  known  -in  early  times.  They  are  the  metals 
universally  used  for  coinage  and  for  ornamental  purposes.  They  are 
the  three  best  conductors  of  electricity  (p.  533),  and  each  represents 
the  maximum  of  conductivity  in  the  periodic  series  to  which  it  belongs. 
In  malleability  and  ductility  silver  is  intermediate  between  gold  and 
copper  (p.  531),  but  in  electrical  conductivity  it  exceeds  both. 

The  Chemical  Relations  of  the  Copper  Family.  —  Copper  (Cu, 
at.  wt.  63.6),  silver  (Ag,  at.  wt.  107.93),  and  gold  (An,  at.  wt.  197.2), 
occupy  the  right  side  in  the  second  column  of  the  table  of  the  periodic 
system  (p.  411),  and  the  chemical  relations  (p.  226)  of  these  elements 
are  in  many  ways  in  sharp  contrast  to  those  of  the  alkali  metals,  their 
neighbors,  on  the  left  side  : 


ALKALI  METALS. 

All  univalent  and  give  but  one  series 
of  compounds.  Halides  all  soluble 
in  water. 

Very  active ;  rapidly  oxidized  by  air ; 
displace  all  other  metals  from  com- 
bination (E.  M.  series,  p.  362). 

Oxides  and  hydroxides  strongly  basic, 
and  halides  not  hydrolyzed  (p.  534). 


COPPEB,  SILVER,  GOLD. 

Cu1  and  Cu11 :  two  series.  Ag1 :  one 
series.  Au1  and  Aum  :  two  series. 
Halides  of  univalent  series  insoluble. 

Amongst  least  active  metals;  only 
copper  is  oxidized  by  air  ;  displaced 
by  most  other  metals.  Hence, 
found  free  in  nature  (p.  362). 

Oxides  and  hydroxides  feebly  basic 
(except  Ag20)  ;  halides  hydrolyzed 
(except  Ag-halides).  Hence,  basic 
salts  are  numerous. 

Give  no  com-        Frequently  in  anion,  e.g.,  K.Cu(CN)2, 
K.Ag(CN)2,    K.Au02,    K.Au(CN)2, 
and  in  complex  cation,  e.g., 
Ag(NH3)2.OH  and  Cu(NH3)4.(OH)2. 

On  account  of  their  inactivity  towards  oxygen,  and  their  easy 
recovery  from  combination  by  means  of  heat,  silver  and  gold, 
together  with  the  platinum  family,  are  known  as  the  "noble  metals." 

Univalent  copper  and  gold  resemble  in  some  ways  Hg1  and  Tl1, 

614 


Never  found  in  anion. 
plex  cations. 


COPPER,   SILVER,   GOLD  615 

while  bivalent  copper  resembles  Zn",  Mn11,  Ee",  and  M",  and  trivalent 
gold  resembles  Alm  and  Fem.  This  family  is,  in  fact,  not  homoge- 
neous, and  the  close  relation  which,  amongst  metals,  subsists  between 
valence  and  chemical  properties  makes  comparisons  with  elements  of 
entirely  different  families  often  the  most  suggestive. 

COPPER. 

Chemical  Relations  of  the  Element.  —  Copper  is  the  first  metal- 
lic element  showing  two  valences  which  we  have  encountered.  In  such 
cases  two  more  or  less  complete,  independent  series  of  salts  are  known. 
These  are  here  distinguished  as  cuprous  (univalent)  and  cupric  (biva- 
lent) salts.  The  methods  by  which  a  compound  of  one  series  may  be 
converted  into  the  corresponding  compound  of  the  other  series  should 
be  noted. 

The  chief  cuprous  compounds  are  Cu/),  CuCl,  CuBr,  Cul,  CuCN, 
CugS.  There  are  no  cuprous  salts  of  oxygen  acids.  The  cuprous  com- 
pound is  in  each  case  more  stable  (p.  119)  than  the  corresponding  cupric 
compound,  and  is  formed  from  it  either  by  spontaneous  decomposition,  as 
in  the  cases  of  the  iodide  and  cyanide  (2CuI2  — >  2CuI  +  I2),  or  on  heat- 
ing. The  cuprous  halides  and  cyanide  are  colorless  and  insoluble  in 
water ;  but  the  chloride,  being  easily  oxidized  by  air  to  the  cupric  con- 
dition, quickly  turns  green.  The  ion  Cu*  (monocuprion)  seems  to  be 
colorless. 

The  cupric  compounds  are  more  numerous,  as  they  include  also 
salts  of  oxygen  acids,  like  CuS04,  Cu(NOs)2,  etc.  CuI2  and  Cu(CN)2 
cannot  be  obtained  in  pure  form,  as  they  decompose,  giving  the  cuprous 
salts.  The  anhydrous  salts  are  usually  colorless  or  yellow,  but  the 
ion  Cu"  (dicuprion)  is  blue,  and  so,  therefore,  are  the  aqueous  solu- 
tions of  the  salts.  The  cupric  are  more  familiar  than  the  cuprous 
compounds,  since  cupric  oxide,  sulphate,  and  acetate  are  the  compounds 
of  copper  which  most  frequently  find  employment  in  chemistry  and 
in  the  arts.  All  the  soluble  salts  of  copper  are  poisonous. 

In  electrolyzing  salts  of  copper,  a  given  amount  of  electricity  will 
deposit  twice  as  much  copper  from  a  cuprous  salt  as  from  a  cupric  salt 
(p.  316),  since  monocuprion  carries  only  half  as  great  a  charge,  weight 
for  weight,  as  dicuprion. 

Writers  on  chemistry  frequently  double  (CUjClg,  etc.)  the  formulae 
of  cuprous  salts.  The  molecular  weights  in  organic  solvents  (cf.  p.  292), 
however,  in  many  cases  accord  with  the  simple  formulae.  Some  higher 


616  INORGANIC    CHEMISTRY 

molecular  weights  observed  in  solution  and  the  vapor  density  of  cuprous 
chloride  (6.6,  corresponding  nearly  to  Cu2Cl2)  might  be  regarded  as 
being  due  to  association  (imperfect  dissociation,  cf.  p.  205).  The  for- 
mation of  numerous  double  or  complex  compounds  like  HCl,CuCl 
(=  HCuCLj),  which  may  be  regarded  as  acid  salts,  however,  lends  sup- 
port to  the  view  that  the  formulae  should  be  doubled.  Inasmuch  as  the 
behavior  of  the  salts  is  sufficiently  well  represented  by  the  simple  for- 
mulae, these  are  here  used  throughout  (see  under  Silver,  p.  626). 

Occurrence.  —  Copper  is  found  free  in  the  Lake  Superior  region, 
in  China,  and  in  Japan.  The  sulphides,  copper  pyrites  CuFeS2  and 
chalcocite  Cu2S,  are  worked  in  Montana,  in  southwest  England,  in 
Spain,  and  in  Germany.  Malachite,  Cu2(OH)2CO3  (  =  Cu(OH)2,CuCO3), 
and  azurite,  Cu8(OH)2(C03)2(=  Cu(OH)2,2CuC08),  both  basic  carbon- 
ates, are  mined  in  Siberia  and  elsewhere.  Cuprite  or  ruby  copper 
CujjO  is  also  an  important  ore.  The  name  of  the  element  comes  from 
the  fact  that  in  ancient  times  copper  mines,  long  since  worked  out,  ex- 
isted in  Cyprus.  The  element  is  found  in  the  feathers  of  some  birds, 
in  the  hsemocyanin  of  the  blood  of  the  cuttle-fish,  which  is  blue  when 
arterial  and  colorless  when  venous,  and  elsewhere  in  living  organisms. 

Extraction  from  Ores.  —  For  isolating  native  copper  it  is  only 
necessary  to  separate  the  metal,  by  grinding  and  washing,  from  the 
rock  through  which  it  ramifies,  and  to  melt  the  almost  pure  powder  of 
copper  with  a  flux  (p.  540).  The  carbonate  and  oxide  ores  require 
coal,  in  addition,  for  the  removal  of  the  oxygen. 

The  liberation  of  copper  from  the  sulphide  ores  is  difficult,  and  often 
involves  very  elaborate  schemes  of  treatment.  This  arises  from  two 
causes.  Other  metals,  such  as  iron  and  zinc,  unite  with  oxygen  more 
readily  than  with  sulphur,  and  hence  the  sulphides  are  easily  converted 
into  oxides  by  roasting  with  free  access  of  air ;  with  copper  it  is  just  the 
reverse.  Thus,  even  great  excess  of  air  and  repeated  roasting  produce 
only  a  gradual  diminution  in  the  amount  of  sulphur  in  the  mass.  Then, 
many  copper  ores  contain  a  large  amount  of  the  sulphides  of  iron,  and 
these  have  to  be  removed  by  conversion  into  oxide  (by  roasting)  and 
then  into  silicate  (with  sand).  The  silicate  forms  a  flux,  and  separates 
itself  from  the  molten  mixture  of  copper  and  copper  sulphide 
("matte").  In  Montana  it  is  found  possible  to  abbreviate  the  treat- 
ment. The  ore  is  first  roasted  until  partially  oxidized.  It  is  then 
melted  in  a  cupola  or  a  reverberating  furnace,  and  placed  in  large  iron 


COPPER,    SILVER,    GOLD  617 

vessels  like  Bessemer  converters  (q.v.)  provided  with  a  lining  rich  in 
silica.  A  blast  of  air  mixed  with  sand  is  now  blown  through  the 
mass.  The  iron  is  completely  oxidized  to  FeO  and  made  into  silicate, 
the  sulphur  escapes  as  sulphur  dioxide,  and  arsenic  and  lead  are  like- 
wise removed  by  this  treatment.  The  silicate  of  iron  floats  as  a  slag 
upon  the  copper  when  the  contents  of  the  converter  are  poured  out. 
The  resulting  copper  is  pure  enough  to  be  cast  in  large  plates  and 
purified  by  electrolysis  (see  below). 

Wet  processes  are  used  for  poor  ores.  In  one  of  these  the  ore  is 
roasted  with  salt.  The  copper  is  thus  converted  into  cupric  chloride, 
and  can  be  dissolved  away  from  the  oxide  of  iron  and  other  materials 
by  means  of  water.  Any  traces  of  silver  which  may  have  been  present 
pass  also  into  solution  (as  AgCl),  and  are  precipitated  by  addition  of 
potassium  iodide.  The  copper  is  then  displaced  by  means  of  scrap 
iron,  and  forms  a  brown  sludge  (Cu"  -f  Fef  — »  Fe"  +  Cu  J). 

The  properties  of  copper  are  seriously  affected  by  small  amounts  of 
impurities,  such  as  cuprous  oxide  or  sulphide,  which  are  soluble  in  the 
molten  metal.  Then,  too,  the  difficulties  in  the  way  of  its  preparation 
are  aggravated  by  the  high  standard  of  purity  demanded  in  order  that  it 
may  have  the  maximum  tenacity,  ductility,  and  conductivity  which  its 
various  applications  require.  Hence  a  large  proportion  of  the  copper 
on  the  market  is  purified  by  electrolysis,  a  method  which  meets  the 
case  by  giving  a  copper  in  which  foreign  materials  can  be  shown  to  be 
present  only  by  the  most  refined  tests.  In  this  method  of  refining  the 
metal,  thin  sheets  of  copper,  coated  with  graphite  to  permit  easy  re- 
moval of  the  deposit,  form  the  cathodes,  and  thick  plates  of  copper  the 
anodes.  These  are  suspended  alternately  and  close  together  in  large 
troughs  filled  with  cupric  sulphate  solution.  The  cathodes  are  all 
connected  with  the  negative  wire  of  the  dynamo,  and  the  anodes  with 
the  positive  one.  The  Cu"  is  attracted  to  the  cathodes  and  is  deposited 
upon  them.  The  SO/'  migrates  towards  the  anodes,  where  copper 
from  the  thick  plate  becomes  ionized  in  equivalent  amount.  The  stock 
of  cupric  sulphate  thus  remains  the  same,  and  the  practical  effect  of 
the  electrolysis  is  to  carry  copper  across  from  one  plate  to  the  other 
(cf.  p.  325).  The  cathodes  are  removed  from  time  to  time,  and  the 
deposit  of  copper  is  stripped  from  their  surface.  Fresh  anodes  are 
substituted  when  the  old  ones  are  eaten  away.  Since  there  is  no  polar- 
ization, a  current  of  less  than  0.5  volt  suffices.  The  copper  is  deposited 
in  pure  form,  although  an  impure  anode  is  employed,  because  metals 
like  gold,  silver,  and  antimony,  which  succeed  copper  in  the  electro- 


618  INORGANIC   CHEMISTRY 

motive  series  (p.  362),  and  compounds  like  cuprous  sulphide,  do  not 
become  ionized.  They  fall  to  the  bottom  of  the  tank  as  a  fine  powder. 
Similarly,  metals  like  zinc,  which  displace  copper,  although  they  are 
dissolved,  are  not  again  deposited.  This  will  be  understood  when  it  is 
considered  that,  if  they  were  to  be  deposited,  they  would  displace 
copper  from  the  solution  and  dissolve.  Since  the  copper,  as  deposited 
on  the  cathodes,  is  crystalline  and  porous,  it  is  afterwards  melted  and 
cast  into  blocks  or  bars. 

In  1900  the  United  States  produced  60  per  cent  of  the  world's 
copper,  Great  Britain  16  per  cent,  and  Germany  6  per  cent.  The  pro- 
portions of  the  whole  consumed  in  each  of  these  countries  were  33,  22, 
and  22  per  cent,  respectively. 

Physical  Properties.  —  Copper  is  red  by  reflected  and  greenish  by 
transmitted  light.  Native  copper  shows  crystals  of  the  regular  system 
(p.  530).  It  melts  at  1057°,  and  therefore  much  more  easily  than  pure 
iron  (1800°).  When  steel  draw-plates  are  used  it  can  be  drawn  into 
wire  with  a  diameter  of  only '2  mm.,  and  by  means  of  plates  provided 
with  perforated  diamonds  the  diameter  of  the  wire  can  be  reduced  to 
.03  mm.  (1  kilometer  weighs  only  7  g.).  The  metal  after  drawing  is 
more  tenacious  but  conducts  electricity  less  well. 

Chemical  Properties.  —  In  dry  oxygen,  copper  does  not  rust. 
In  moist  oxygen  a  thin  film  of  cuprous  oxide  is  formed,  and  in  ordinary 
air  a  green  basic  carbonate  (not  verdigris,  q.v.).  It  does  not  decom- 
pose water  at  any  temperature  or  displace  hydrogen  from  dilute  acids 
(p.  362).  On  the  other  hand,  hydrogen,  absorbed  in  platinum  or  even 
in  charcoal,  liberates  copper  (Cu  **  -+-  H2  — -»  Cu  -f  2H*)  when  immersed 
in  solutions  of  copper  salts.  The  metal  attacks  oxygen  acids  (pp. 
379,  446),  however.  Again,  acids  like  hydrochloric  acid,  in  conjunc- 
tion with  oxygen  from  the  air,  do  act  slowly  upon  copper  (2Cu  -f-  4HC1 
-f  O2  — >  2CuCl2  -f  2H20).  This  sort  of  simultaneous  action  of  two 
agents  is  frequently  used,  as  in  making  silicon  tetrachloride  (p.  520). 
In  a  similar  way  sea-water  and  air  slowly  corrode  the  copper  sheath- 
ings  of  ships,  giving  a  basic  chloride,  Cu4(OH)6Cl2,  H20(=  3Cu(OH)2, 
Cu012,H20),  which  is  found  in  nature  as  atakarnite. 

On  account  of  its  resistance  to  the  action  of  acids,  copper  is  used 
for  many  kinds  of  vessels,  for  covering  roofs  and  ships7  bottoms,  and 
for  coins.  It  furnishes  also  electrotype  reproductions  of  medals,  of 
engraved  plates,  of  type,  etc.  For  this  purpose  a  cast  of  the  object 


COPPER,    SILVER,   GOLD  619 

is  first  made  in  gutta  percha,  plaster  of  paris,  or  wax.     This  is  then 
'  coated  with  graphite  to  give  it  a  conducting  surface,  and  receives  an 
electrolytic  deposit  of  copper.      Great  quantities  of  the  metal  are  used 
in  electrical  plants  and  appliances. 

Alloys.  —  The  qualities  of  copper  are  modified  for  special  purposes 
by  alloying  it  with  other  metals.  Brass  contains  18-40  per  cent  of 
zinc,  and  melts  at  a  lower  temperature  (p.  532)  than  does  copper.  A 
variety  with  little  zinc  is  beaten  into  thin  sheets,  giving  Dutch-metal 
("  gold-leaf ").  Bronze  contains  3-8  per  cent  of  tin,  11  or  more  per 
cent  of  zinc,  and  some  lead.  It  was  used  for  making  weapons  and 
tools  before  means  of  hardening  iron  were  known,  and  later,  on  account 
of  its  fusibility,  continued  to  be  employed  for  castings  until  displaced 
largely  by  cast-iron  (discovered  in  the  eighteenth  century).  For  works 
of  art  it  is  preferred  to  copper  because  of  its  fusibility,  its  color,  and 
its  more  rapid  acquirement  of  a  much  prized  "  patina  "  due  to  surface 
corrosion.  Artificial  "  bronzing "  of  brass  is  effected  by  applying  a 
solution  of  arsenious  oxide  in  hydrochloric  acid  (AsCl8).  The  zinc 
displaces  some  arsenic,  which  combines  with  the  copper.  Brass  instru- 
ments are  "  bronzed  "  by  means  of  a  dilute  solution  of  chloroplatinic 
acid  (q-v.),  from  which  the  zinc  displaces  platinum.  Gun-metal  con- 
tains 10  per  cent,  and  bell-metal  25  per  cent  of  tin.  German  silver 
contains  19-44  per  cent  of  zinc  and  6-22  per  cent  of  nickel,  and  shows 
none  of  the  color  of  copper.  Aluminium-bronze  contains  5-10  per  cent 
of  aluminium,  and  resembles  gold  in  color.  When  it  contains  some 
iron  it  can  be  worked  at  a  red  heat,  but  not  welded.  Silicon-bronze 
contains  not  more  than  5  per  cent  of  silicon,  and  is  made  by  adding 
silicide  of  copper  (made  in  the  electric  furnace,  p.  457)  to  copper.  It 
has  usually  only  60  per  cent  of  the  conductivity  of  pure  copper,  but  is 
nearly  twice  as  tenacious,  and  is  used  for  telephone  and  over-head 
electric  wires.  Phosphor-bronze  contains  copper  and  tin  (100  :  9)  with 
J-l  part  of  phosphorus,  and  is  employed  for  certain  parts  of  machines. 
Ships'  propellers  are  made  of  manganese-bronze  (30  per  cent  manganese). 

Cupric  Chloride.  —  This  compound  is  made  by  union  of  copper 
and  chlorine,  by  treating  the  hydrate  or  carbonate  with  hydrochloric 
acid,  or  by  heating  copper  with  hydrochloric  acid  and  some  nitric  acid 
the  latter  being  used  simply  as  an  oxidizing  agent  (Cu  +  2HC1  +  0  — > 
CuCLj  +  H20).  The  blue  crystals  of  a  hydrate,  CuCLj,2H2O,  are  de- 
posited by  the  solution.  The  anhydrous  salt  is  yellow.  Dilute  solu- 


620  INORGANIC   CHEMISTRY 

tions  are  blue,  the  color  of  dicuprion,  but  concentrated  solutions  are 
green  on  account  of  the  presence  of  the  yellow  molecules  (p.  335). 
The  aqueous  solution  is  acid  in  reaction  (p.  344).  When  excess  of 
ammonium  hydroxide  is  added  to  the  solution,  the  basic  chloride,  cupric 
oxychloride  Cu4(OH)6CLj  (p.  618),  which  is  at  first  precipitated,  redis- 
solves,  and  a  deep-blue  solution  is  obtained.  This  on  evaporation  yields 
deep-blue  crystals  of  hydrated  ammonio-cupric  chloride  Cu(NH8)4. 
CLj,  ELjO.  The  deep-blue  color  of  the  solution,  which  is  given  by  all 
cupric  salts,  is  that  of  the  Cu(NH8)4"  ion.  The  dry  salt  also  absorbs 
ammonia,  giving  CuCLj,  6NH3.  This  and  the  preceding  compound  have 
an  appreciable  tension  of  ammonia,  and  when  warmed  leave  CuCLj,  2NH3. 
Reduction  of  pressure  or  rise  of  temperature  results  in  the  final  loss 
of  all  the  ammonia  (cf.  p.  122). 

Cuprous  Chloride.  —  This  salt  is  formed  when  cupric  chloride  is 
heated  (2CuCLj  <=±  2CuCl  +  CLj).  It  may  be  made  by  adding  hydro- 
chloric acid  to  cupric  chloride  solution,  and  boiling  the  mixture  with 
copper  turnings : 

CuCLj  +  Cu  ->  2CuCl  or  Cu"  +  Cu  ->  2Cu*. 

The  solution  contains  compounds  of  cuprous  chloride  with  hydrogen 
chloride  HCl,CuCl  or  HCuClg  and  H2CuClg,  which  are  decomposed  when 
water  is  added.  The  cuprous  chloride  is  insoluble  in  water,  and  forms 
a  white  crystalline  precipitate. 

Cuprous  chloride  is  hydrolyzed  quickly  by  hot  water,  giving, 
finally,  red,  hydrated  cuprous  oxide,  Cu20.  When  dry  it  is  not 
affected  by  light,  but  in  the  moist  state  becomes  violet  and,  finally, 
nearly  black.  The  action  is  said  to  be  2CuCl  ->  CuCl2  +  Cu.  In 
moist  air  it  turns  green,  and  is  oxidized  to  cupric  oxychloride  (p.  618). 
It  is  dissolved  by  hydrochloric  acid,  giving  the  colorless  complex  acids 
HCuCLj  and  H2CuCl3.  The  solution  is  oxidized  by  the  air,  turning 
first  brown  and  then  green,  and  finally  depositing  the  cupric  oxychlo- 
ride. Cuprous  chloride  also  dissolves  in  ammonium  hydroxide,  giving 
Cu(NH8)2.Cl,  the  ion  Cu(NH3)2*  being  colorless.  The  solution  is 
quickly  oxidized  by  the  air,  turns  deep-blue,  and  then  contains 
Cu(NH3)4".  The  solution  of  cuprous  chloride  in  hydrochloric  acid  is 
used  for  absorbing  carbon  monoxide  from  gaseous  mixtures.  A  crys- 
talline compound  (CuCl,  CO,  2H2O  ?)  has  been  isolated  from  the  solu- 
tion. 


COPPER,    SILVER,   GOLD  621 

The  Bromides  and  Iodides  of  Copper.  —  By  treatment  of  cu- 
pric  oxide  with  hydrobromic  acid  and  slow  evaporation  of  the  solution, 
jet-black  crystals  of  anhydrous  cupric  bromide  (CuBr2)  are  obtained. 
A  concentrated  aqueous  solution  is  deep-brown  in  color,  and  the  grad- 
ual ionization  of  the  molecules  as  the  solution  is  diluted  is  well  shown  by 
this  salt  (p.  335).  The  ionization  is  here  accompanied  by  evolution  of 
heat  (p.  330),  as  it  is  also  in  the  cases  of  cupric  chloride  and  cupric  sul- 
phate, and  in  the  ionized  condition  the  substances  contain  less  available 
energy  than  in  the  molecular.  In  these  cases,  therefore,  when  the  tem- 
perature is  raised  the  ionization  diminishes  (p.  260).  It  will  be  remem- 
bered that  ordinary  thermal  dissociation  usually  increases  with  rise 
in  temperature,  and  is  then  accompanied  by  absorption  of  heat.  Ionic 
dissociation,  however,  may  be  either  endothermal  or  exothermal. 

When  cupric  bromide  is  heated,  bromine  is  given  off,  and  cuprous 
bromide  CuBr  remains. 

Cupric  iodide  appears  to  be  unstable  at  ordinary  temperatures. 
When  a  soluble  iodide  is  added  to  a  cupric  salt,  a  white  precipitate  of 
cuprous  iodide  and  free  iodine  are  obtained  : 


The  iodine  may  be  dissolved  in  excess  of  the  soluble  iodide  (p.  235),  or 
reduced  to  hydrogen  iodide  with  sulphurous  acid  (p.  394). 

The  Solution  of  Insoluble  Salts  when  Complex  Ions  are 
Formed.  —  The  solution  of  an  insoluble  salt  like  cuprous  chloride  by 
hydrochloric  acid  or  ammonium  hydroxide  is  typical  of  a  great  variety 
of  actions  of  which  we  here  meet  one  of  the  first  examples  (cf.  p.  363). 

Since  a  salt  is  normally  less  soluble  in  an  acid  having  the  same 
anion  (p.  585),  the  dissolving  of  cuprous  chloride  in  hydrochloric  acid 
requires  a  special  explanation,  namely,  the  fact  that  here  two  complex 
acids  H.CuClj  and  H2.CuCl3  are  formed.  The  chloridion  of  the  hy- 
drogen chloride  must  indeed  tend  to  repress  the  ionization  of  the  dis- 
solved part  of  the  cuprous  chloride,  so  that  a  smaller  concentration  of 
Cu*  remains.  But  the  complex  negative  ion  CuCL/  (or  CuCl8")  which  is 
formed,  is  very  little  dissociated,  and  gives  a  still  smaller  concentration 
of  Cu*  (CuCl/  <=»  Cu*  -f  2C1').  Thus  this  complex  ion  is  formed  at  the 
expense  of  the  Cu"  of  the  insoluble  cuprous  chloride,  and  the  latter 
goes  ^nto  solution  progressively  in  the  effort  to  restore  the  balance  : 

CuCl  (solid)  ±5  CuCl  (diss'd)  <=±  Cl'  +  Cu' 
2HC1  =±2 


622  INORGANIC   CHEMISTRY 

The  same  exact  laws  of  equilibrium  used  in  discussing  the  dissolving 
of  salts  by  acids  with  a  different  anion  (p.  600)  may  be  applied  to  the 
whole  procedure. 

Similar  behavior  is  shown  by  the  cyanides  of  copper,  silver,  iron, 
etc.  (q-v.),  of  which  many  complex  compounds  are  known. 

The  dissolving  of  cuprous  chloride  by  the  free  ammonia  of  ammo- 
nium hydroxide  is  explained  in  the  same  way.  The  only  difference 
is  that  here  the  copper  is  in  the  complex  positive  ion.  The  ion 
Cu(NH8)2*  gives  little  Cu* —  less  than  does  cuprous  chloride,  in  spite  of 
the  insolubility  of  the  latter.  Hence  the  salt  passes  into  solution 
until  the  ion-product  [Cu*]  x  [Cl'],  with  continually  increasing  [Cl'], 
reaches  its  normal  value  or  until  the  solid  is  exhausted. 

The  deep-blue  colored  ion  Cu(NH3)4"  given  by  cupric  chloride  and 
other  cupric  salts  is  also  very  little  ionized.  Hence  ammonium  hy- 
droxide dissolves  all  the  insoluble  cupric  compounds  save  only  cupric 
sulphide,  which  is  the  most  insoluble  of  all  —  that  is,  the  one  giving 
the  smallest  concentration  of  dicuprion.  Conversely,  the  sulphide  is 
the  only  insoluble  compound  of  copper  which  can  be  precipitated  from 
ammoniacal  solution.  Zinc  and  other  more  active  metals,  however, 
slowly  precipitate  metallic  copper,  thereby  showing  that  some  dicup- 
rion is  present. 

Cuprous  Oxide. — This  oxide  is  red  in  color,  and  natural  specimens 
show  octahedral  forms.  It  is  produced  by  oxidation  of  finely  divided 
copper  at  a  gentle  heat,  or  by  the  addition  of  bases  to  cuprous  chloride, 
and  is  best  made  by  the  action  of  glucose  on  cupric  hydroxide.  The 
latter  is  reduced  by  the  former,  and  the  resulting  hydrated  cuprous 
oxide  forms  a  pale-brown  precipitate  which  quickly  becomes  bright 
red.  The  simple  hydrate,  CuOH,  is  unknown,  but  the  above  mentioned 
precipitate  has  approximately  the  composition  4Cu20,H20,  and  yields 
Cu20  when  heated. 

Cuprous  oxide  is  acted  upon  by  hydrochloric  acid,  giving  cuprous 
chloride,  or  rather  HCuCLj.  It  also  dissolves  in  ammonium  hydroxide, 
giving,  probably,  Cu(NH8)2.OH,  which  is  colorless.  With  dilute 
oxygen  acids  part  of  it  is  oxidized,  giving  the  cupric  salt,  and  part  is 
reduced  to  metallic  copper : 

Cu20  +  H2S04  -*  CuS04  +  Cu  +  H20. 

Cupric  Oxide  and  Hydroxide.  —  Cupric  oxide  CuO  is  a  black 
substance  formed  by  heating  copper  in  a  stream  of  oxygen  or  by  igniting 


COPPER,    SILVER,   GOLD  623 

the  nitrate,  carbonate,  or  hydroxide.  It  absorbs  moisture  from  the  air, 
3,1  though  it  is  not  soluble  in  water.  When  heated  strongly  it  loses  some 
Dxygen,  and  is  partly  reduced  to  cuprous  oxide.  Its  chief  use  is  in  the 
analysis  of  compounds  of  carbon.  When  heated  with  the  latter,  it 
oxidizes  the  hydrogen  to  water,  and  the  carbon  to  carbon  dioxide.  The 
operation  is  performed  in  a  tube  through  which  passes  a  stream  of 
oxygen,  and  the  products  are  caught  in  glass,  vessels  containing  calcium 
chloride  and  potassium  hydroxide,  respectively. 

Cupric  hydroxide  is  precipitated  as  a  gelatinous  substance  by 
addition  of  sodium  or  potassium  hydroxide  to  a  solution  of  a  cupric 
salt  (Cu-  +  20H' — >  Cu(OH)2).  When  the  mixture  is  boiled,  the 
hydroxide  loses  water  and  forms  a  black  hydrated  cupric  oxide 
(Cu(OH)2, 2CuO  ?).  The  hydroxide  is  soluble  in  ammonium  hydroxide, 
with  formation  of  the  compound  Cu(NH3)4.(OH)2,  which  imparts  a 
deep-blue  color  to  the  solution.  Various  forms  of  cellulose,  such  as 
filter  paper  and  cotton,  dissolve  in  this  solution,  and  are  reprecipitated 
when  the  ammonium  hydroxide  is  neutralized  with  acids.  Cupric 
hydroxide  dissolves  also  in  a  solution  of  sodium  tartrate  (Na2.C4H2O4 
(OH)2),  giving  a  deep-blue  liquid  (practically  "Fehling's  solution"), 
in  this  action,  it  enters  into  the  negative  ion,  as  is  shown  by  electrolysis, 
interacting  apparently  with  the  hydroxyl  groups  of  the  tartranion. 
In  this  condition  it  is  reduced  to  cuprous  oxide  by  sugars  (p.  622) 
with  especial  ease,  and  is  in  this  form  used  as  a  test  for  them. 


Cupric  Nitrate.  —  The  nitrate  is  made  by  treating  cupric  oxide 
or  copper  with  nitric  acid  (p.  446),  and  is  obtained  from  the  solution 
as  a  deliquescent,  crystalline  hydrate.  The  hexahydrate  is  secured 
at  temperatures  below  24.5°,  its  transition  point  (p.  594),  and  the  tri- 
hydrate  from  24.5°  up  to  114.5°  (its  transition  point;  see  under  Man- 
ganous  sulphate).  When  dehydrated  at  65°  the  salt  is  partly 
hydrolyzed,  and  a  basic  nitrate  Cu4(OH)6(N08)2  remains. 


Carbonate  of  Copper.  —  No  normal  carbonate  (CuC03)  can  be 
obtained.  A  basic  carbonate  (malachite)  is  found  in  nature,  and  is 
precipitated  by  adding  soluble  carbonates  to  cupric  salts  : 


2CuS04  +  2X^00,  +  H20  ->  Cu2(OH)2C08  +  2NajS04  +  C02. 
Presumably,  the  carbonate,  if  formed,  would  be  hydrolyzed  by  water. 


624  INORGANIC   CHEMISTRY 

Cyanides  of  Copper.  —  When  potassium  cyanide  is  added  k>  a 
solution  of  a  cupric  salt,  cupric  cyanide  is  precipitated.  This  is  not 
stable,  however,  and  gives  off  cyanogen,  leaving  cuprous  cyanide  : 

2Cu(CN)2  ->  2CuCN  +  C2N2. 

Cuprous  cyanide  is  insoluble  in  water,  but  interacts  with  an  excess  of 
potassium  cyanide  solution,  producing  a  colorless  liquid,  from  which 
KCNjCuCN,  or  K.Cu(CN)2,  potassium  cuprocyanide,  may  be  obtained 
in  colorless  crystals.  The  complex  anion  Cu(CN)2'  is  so  little  ionized 
to  Cu"  and  2CN'  that  all  insoluble  copper  compounds,  including  cupric 
sulphide,  are  dissolved  by  potassium  cyanide  ;  and  none  of  them  can 
be  precipitated  from  the  solution.  Zinc  is  actually  unable  to  displace 
copper  from  such  a  solution.  The  cause  of  the  solution  of  the  salts 
is  the  same  as  when  the  complex  ions  Cu(NH8)2',  Cu(NH8)4",  and 
are  formed  (p.  621). 


Cupric  Acetate.  —  By  the  oxidation  of  plates  of  copper,  sep- 
arated by  cloths  saturated  with  acetic  acid  (vinegar),  a  basic  acetate 
of  copper  (verdigris)  is  obtained  : 

6Cu  +  8HC2H802  +  302  _>  2Cu8(OH)2(C2H802)4  +  2H2O. 

It  is  used  in  manufacturing  green  paint,  is  insoluble  in  water,  and  is  un- 
affected by  light.  It  dissolves  in  acetic  acid,  and  green  crystals  of  the 
normal  acetate  Cu(C2H802)2,H20  are  obtained  from  the  solution.  The 
basic  acetate  is  used  in  preparing  paris  green.  A  hot  solution  of 
arsenious  acid  (H8As08)  is  mixed  with  a  paste  of  verdigris  and  a  little 
acetic  acid  and  boiled.  A  precipitate  of  paris  green  Cu(C2H802)2,Cu3 
As206,  which  has  a  unique  light-green  color,  is  thrown  down.  On  ac- 
count of  their  poisonous  nature,  this  compound  and  Scheele's  green 
(CuHAs03)  are  little  used  as  pigments.  The  former  is  chiefly  made 
for  use  in  the  extermination  of  potato-beetles  and  other  insects  and 
employment  in  the  destruction  of  parasitic  fungi. 

Cupric  Sulphate.  —  This  salt  is  obtained  by  heating  copper  in  a 
furnace  with  sulphur,  and  admitting  air  to  oxidize  the  cuprous  sul- 
phide. The  mixture  of  cupric  sulphate  and  cupric  oxide  which  is 
formed  is  treated  with  sulphuric  acid.  The  salt  is  also  made  by  allow- 
ing dilute  sulphuric  acid  to  trickle  over  granulated  copper  while  air 
has  free  access  to  the  material  (2Cu  -f  2HaSO,  +  Oa  -»  2CuS04  + 


COPPER,    SILVER,    GOLD  625 

2H20).  When  concentrated  and  at  a  high  temperature,  sulphuric 
acid  will  itself  act  as  the  oxidizing  agent  (cf.  p.  379). 

Cupric  sulphate  crystallizes  as  pentahydrate  CuS04,5H2O  in  blue 
asymmetric  crystals  (Fig.  41,  p.  120),  and  in  this  form  is  called  blue-stone 
or  blue  vitriol.  The  dissociation  of  this  hydrate  has  been  discussed 
on  page  122.  The  aqueous  solution  has  an  acid  reaction  (p.  344). 
The  anhydrous  salt  is  white,  and  can  be  crystallized  in  thin  needles 
(rhombic  system?)  from  solution  in  hot,  concentrated  sulphuric  acid 
(cf.  pp.  120-123).  Cupric  sulphate  is  employed  for  making  other  com- 
pounds of  copper,  in  copper-plating  (p.  618),  in  batteries,  as  a  mordant 
in  dyeing  (q.v.)  and  calico-printing,  and,  as  a  germicide  and  insecti- 
cide, for  spraying  plants. 

When  ammonium  hydroxide  is  added  to  cupric  sulphate  solution, 
a  pale-green  basic  salt  (Cu4(OH)6S04  ?)  is  first  precipitated.  With 
excess  of  the  hydroxide  the  blue  Cu(NH3)4"*  ion  (p.  620)  is  formed,  and 
crystals  of  ammonio-cupric  sulphate  Cu(NH3)4.  S04,H2O  can  be  obtained 
from  the  solution.  This  compound  easily  loses  water  and  ammonia 
(by [stages),  leaving  successively  CuSO4, 2NH3  and  CuS04,  NH3.  Cupric 
sulphate  also  combines  with  potassium  and  ammonium  sulphates, 
giving  double  salts  of  the  form  CuS04,K2S04,6H20,  which  are  de- 
posited in  large,  monosymmetric  crystals  from  the  mixed  solutions 
(see  Zinc  sulphate). 

The  Sulphides  of  Copper.  —  Cuprous  sulphide  Cu2S-  occurs  in 
nature  in  rhombic  crystals  of  a  gray,  metallic  appearance.  It  is  made 
by  heating  cupric  sulphide,  a  stream  of  hydrogen  gas  being  used  to 
assist  the  removal  of  the  excess  of  sulphur. 

Cupric  sulphide  is  deposited  as  a  black  precipitate  when  hydrogen 
sulphide  is  led  through  a  solution  of  a  cupric  salt.  Made  in  this  way, 
it  is  always  partly  decomposed  into  Cu2S  +  S.  By  cautiously  treating 
copper  with  excess  of  sulphur  at  114°  it  may  be  obtained  as  a  blue 
crystalline  solid.  At  higher  temperatures  it  gives  off  sulphur. 

Analytical  Reactions  of  Compounds  of  Copper.  —  The  ion  of 

ordinary  cupric  salts,  dicuprion  Cu",  is  blue,  and  that  of  cuprous 
salts,  monocuprion  Cu*,  is  colorless.  Cuprous  solutions,  however,  are 
easily  oxidized  by  the  air  and  become  blue.  In  solutions  containing 
dicuprion,  hydrogen  sulphide  precipitates  cupric  sulphide,  even  in 
presence  of  acids  (p.  600).  Bases  throw  down  the  blue  hydroxide, 
and  carbonates  precipitate  a  green  basic  salt  (p.  623)  Potassium  fer- 


626  INORGANIC   CHEMISTRY 

rocyanide  gives  the  brown,  gelatinous  cupric  ferrocyanide  (2Cu.S04  + 
K4.Fe(CN)6  <=±  Cu2.Fe(CN)6  J  +  2K2S04).  A  very  characteristic  test  is 
the  formation  of  the  deep-blue  Cu(NH8)4"  ion  with  excess  of  ammo- 
nium hydroxide.  This  solution  itself  gives  a  precipitate  with  hydro- 
gen sulphide  only.  Solutions  of  complex  cuprous  and  cupric  cyanides 
such  as  K.Cu(CN)2  and  K2.Cu(CN)4  are  colorless,  and  do  not  respond  to 
any  of  the  above  tests.  With  microcosmic  salt  or  borax  (pp.  468,  528), 
copper  compounds  form  a  bead  which  is  green  in  the  oxidizing  part  of 
the  flame  and  becomes  red  and  opaque  (liberation  of  copper)  in  the 
reducing  flame. 

SILVER. 

Chemical  Relations  of  the  Element.  —  This  element  presents  a 
curious  assortment  of  chemical  properties.  It  differs  from  copper  in 
having  a  strongly  basic  oxide,  in  giving  salts  with  active  acids  which 
are  not  hydrolyzed  by  water,  and  in  forming  neutral  rather  than  basic 
salts.  In  these  respects  it  approaches  the  metals  of  the  alkalies  and 
alkaline  earths.  It  resembles  copper  in  entering  into  complex  com- 
pounds, and  in  giving  insoluble  halides  like  the  cuprous  halides.  It 
differs  from  both  copper  and  the  metals  of  the  alkalies,  and  resembles 
gold  and  platinum,  in  that  its  oxide  is  easily  decomposed  by  heat, 
with  formation  of  the  free  metal,  and  in  the  low  position  it  occupies  in 
the  electromotive  series  and  the  consequent  slight  chemical  activity  of 
the  free  metal. 

The  salts  are  always  represented  by  the  simplest  formula,  AgCl, 
etc.,  although  in  organic  solvents  greater  tendencies  to  polymerization 
are  observed  than  in  the  case  of  the  cuprous  compounds  (p.  615). 

Occurrence.  — Native  silver,  sometimes  found  in  large  masses,  al- 
though more  usually  scattered  through  a  rocky  matrix,  contains  vary- 
ing amounts  of  gold  and  copper.  Native  copper  always  contains 
dissolved  silver.  Sulphide  of  silver  (Ag2S)  occurs  alone  and  dissolved 
in  galenite  (PbS),  with  which  it  is  isomorphous.  Smaller  amounts  of 
the  metal  are  obtained  from  pyrargyrite  Ag3SbS3  and  proustite  Ag8AsS3, 
which  are  silver  sulphantimonite  and  sulpharsenite  respectively,  and 
from  horn-silver  AgCl. 

Metallurgy.  — The  silver  contained  free,  or  as  sulphide,  in  ores 
of  copper  and  lead,  is  found  in  the  free  state  dissolved  in  the  metals 
extracted  from  these  ores,  and  is  secured  by  refining  them.  In  the  elec- 
trolytic refining  of  copper,  silver  is  obtained  from  the  mud  deposited  in 


COPPER,   SILVER,   GOLD  627 

the  baths  (p.  617).  The  proportion  present  in  lead  is  usually  small. 
Formerly  the  Pattinson  desilverizing  process  was  largely  employed. 
In  it  the  metallic  lead  is  melted  in  iron  vessels,  and  the  crystals  of 
lead,  deposited  as  the  metal  slowly  loses  heat,  are  raked  out.  These 
consist  at  first  of  pure  lead  (cf.  p.  294).  When  the  remaining  liquid 
becomes  saturated  with  silver  it  begins  to  deposit  lead  and  silver  to- 
gether. At  this  point  the  residue  is  plac.ed  in  a  hollow,  lined  with 
bone-ash,  forming  part  of  a  reverberatory  furnace  (Fig.  100,  p.  572),  and 
heated  strongly  while  a  blast  of  air  passes  over  its  surface.  In  this 
process,  called  "  cupellation,"  the  lead  is  converted  into  litharge  (PbO), 
which,  driven  by  the  air,  flows  in  molten  condition  over  the  edge  of 
the  cupel.  When  the  last  trace  of  lead  is  gone,  the  shining  surface  of 
the  pure  silver  "flashes"  into  view  (cf.  p.  547).  Parke's  process, 
which  has  superseded  the  above,  takes  advantage  of  the  fact  that  mol- 
ten zinc  and  lead  are  practically  insoluble  in  one  another,  while  silver 
is  much  more  soluble  in  zinc  than  in  lead.  Lead  dissolves  1.6  per 
3ent  of  zinc,  and  zinc  1.2  per  cent  of  lead.  The  principle  is  the  same 
is  in  the  removal  of  iodine  from  water  by  ether  (p.  155).  The  lead  is 
melted  and  thoroughly  mixed  by  machinery  with  a  small  proportion  of 
zinc.  After  a  short  time  the  zinc  floats  to  the  top,  carrying  with  it  in 
solution  almost  all  of  the  silver,  and  solidifies  at  a  temperature  at 
which  the  lead  is  still  molten.  The  zinc-silver  alloy  is  skimmed  off, 
and  heated  moderately  in  a  furnace  to  permit  the  adhering  lead  to 
drain  away.  The  zinc  is  finally  distilled  off  in  clay  retorts,  and  the 
lead  remaining  with  the  silver  is  removed  by  cupellation. 

Ores  of  silver  which  do  not  contain  much  or  any  lead  are  often 
smelted  with  lead  ores,  and  the  product  is  treated  as  described  above,  but 
many  other  processes  are  in  use.  Thus,  sulphide  ores  are  sometimes 
roasted  until  the  iron  and  part  of  the  copper  are  converted  into  oxide 
while  the  rest  of  the  copper  and  all  the  silver  remain  as  sulphate. 
The  metal  is  secured  by  extracting  the  mass  with  water  and  precipitat- 
ing the  silver  by  means  of  copper  (p.  362) .  Some  ores  are  roasted  with 
salt,  and  the  resulting  chloride  of  silver  is  dissolved  out  with  sodium 
thiosulphate,  or  even  strong  brine.  In  Mexico  the  "  patio  "  process  has 
been  in  use  since  1557.  The  sulphide  is  converted  into  chloride  by 
the  action  of  cupric  chloride.  Metallic  mercury  displaces  the  silver 
[AgCl  -+-  Hg  — >  HgCl  -f-  -A-g),  and,  being  present  in  excess,  dissolves 
it.  The  treatment  occupies  several  weeks,  and  much  mercury  is  con- 
sumed. The  amalgam  is  finally  secured  by  "  washing,"  and  the  mercury 
is  separated  from  the  silver  by  distillation. 


628  INORGANIC   CHEMISTRY 

In  1899  the  production  of  silver  in  the  United  States  was  2915  tons, 
in  Mexico  1418  tons,  in  Europe  1138  tons.  During  the  first  half  of 
the  nineteenth  century  the  total  world's  output  averaged  only  643  tons 
per  year.  Up  to  1870  a  gram  of  gold  could  buy  15.5  g.  of  silver.  Now 
that  the  production  has  reached  6000  tons,  the  same  amount  of  gold 
purchases  about  35  g. 

Physical  Properties.  —  Pure  silver  is  almost  perfectly  white.  It 
melts  at  960°.  Its  ductility  is  so  great  that  wires  can  be  drawn  of 
such  fineness  that  2  kilometers  of  the  finest  wire  weigh  only  about  1  g. 
In  the  molten  condition  it  absorbs  mechanically  about  twenty-two 
times  its  own  volume  of  oxygen,  but  gives  up  almost  all  of  this  as 
it  solidifies.  Fantastically  irregular  masses  result  from  the  "  sprout- 
ing "  or  "  spitting "  which  accompanies  the  escape  of  the  gas. 

By  addition  of  ferrous  citrate  to  silver  nitrate,  a  red  solution  and 
lilac  precipitate  of  free  silver  can  be  made.  The  latter,  after  washing 
with  ammonium  nitrate  solution,  gives  a  red  solution  in  water.  Other 
solutions  of  colloidal  (cf.  p.  523)  silver  showing  a  variety  of  colors 
have  been  prepared  by  Gary  Lea.  Such  colloidal  solutions  of  metals 
are  formed  also  by  passing  an  electrical  discharge  between  wires  of  sil- 
ver, gold,  or  platinum  held  under  water. 

Silver  is  alloyed  with  copper  to  render  it  harder.  The  silver  coin- 
age of  the  United  States  and  the  continent  of  Europe  has  a  "  fine- 
ness of  900  "  (900  parts  of  silver  in  1000),  and  that  of  Great  Britain 
925.  Silver  ornaments  have  a  fineness  of  800  or  more.  A  superficial 
layer  of  almost  pure-white  silver  is  produced  by  heating  the  object  in 
the  air  and  dissolving  out  the  cupric  oxide  thus  formed  with  dilute 
sulphuric  acid.  The  surface  of  the  products,  if  not  subsequently  bur- 
nished, is  "frosted.77  "Oxidized  silver"  is  made  by  dipping  objects 
made  of  the  metal  in  a  solution  of  potassium  hydrogen  sulphide, 
whereby  a  thin  film  of  silver  sulphide  is  produced. 

Chemical  Properties.  —  Silver  does  not  combine  with  oxygen, 
either  in  the  cold  or  when  heated.  It  does  not  ordinarily  displace 
hydrogen  from  aqueous  solutions  of  acids,  but  its  tendency  to  form  the 
sulphide  is  so  great  that  it  decomposes  hydrogen  sulphide  and  alkali 
sulphides  (cf.  p.  391).  It  also  displaces  hydrogen  when  boiled  with 
concentrated  hydriodic  acid,  giving  Agl.HI.  Silver  interacts  with 
cold  nitric  acid  and  with  hot,  concentrated  sulphuric  acid,  giving  the 
nitrate  or  sulphate  of  silver  and  oxides  of  nitrogen  or  of  sulphur  (p.  446). 


COPPER,    SILVER,    GOLD  629 

Since  its  hydroxide  has  no  tendency  to  behave  as  an  acid,  alkalies, 
whether  in  solution  or  fused,  have  no  action  upon  silver.  Hence  alka- 
line substances  are  heated  in  vessels  of  this  metal  or  of  iron,  rather 
than  in  vessels  of  platinum  (q.v.),  because  platinum,  is  attacked  by 
alkaline  materials. 

The  Halides  of  Silver.  —  The  chloride,  bromide,  and  iodide  are 
formed  as  curdy  precipitates  when  a  salt  of  silver  is  added  to  a  solu- 
tion containing  the  appropriate  halide  ion.  The  first  is  white,  and 
melts  at  about  457°.  The  second  and  third  are  very  pale-yellow  and 
yellow  respectively.  The  insolubility  in  water,  which  is  very  great, 
increases  in  the  above  order.  The  iodide,  after  melting,  solidifies  and 
forms  quadratic  crystals,  which,  as  they  cool,  pass  at  146°  into  a 
different  physical  variety  (hexagonal)  with  evolution  of  heat  (cf. 
pp.  368,  565). 

When  exposed  to  light,  the  chloride  becomes  first  violet  and  finally 
brown,  chlorine  being  liberated.  The  bromide  and  iodide  behave  simi- 
larly. It  is  believed  that  a  sub-chloride  and  sub-bromide  Ag2Cl  and 
Ag2Br  are  formed  in  the  early  stages  of  the  action  (see  Photography, 
below).  Solid  silver  chloride  absorbs  ammonia,  forming  first  2AgCl, 
3NH3,  and  then  AgCl,3NH3,  the  former  with  a  tension  of  93  mm.,  and 
the  latter  with  a  tension  of  about  one  atmosphere  of  ammonia  at  20° 
(cf.  p.  123).  The  bromide  forms  no  compound  in  this  way,  but  the 
iodide  yields  2AgI,NH8. 

In  consequence  of  the  progressive  insolubility,  a  cold  solution  of  a 
bromide  will  slowly  convert  the  precipitate  of  silver  chloride  into 
bromide,  and  a  soluble  iodide  will  similarly  transform  the  bromide  or 
the  chloride  into  iodide  (cf.  p.  585).  Chlorine  gas,  however,  displaces 
bromine  and  iodine  from  the  dry  compounds  (cf.  p.  361).  This  illus- 
trates well  the  absence  of  any  relation  between  the  electromotive 
series  and  double  decomposition  (p.  362). 

Silver  fluoride  may  be  made  by  treating  the  oxide  or  carbonate 
with  hydrofluoric  acid  (H^  +  Ag20  ->  2AgF  +  H2O).  The  salt  is 
very  soluble  and  deliquescent. 

Complex  Compounds  of  Silver.  —  Silver  chloride  dissolves  easily 
in  excess  of  ammonium  hydroxide,  giving  the  complex  cation  Ag(NH8)2*. 
Under  certain  conditions  octahedral  crystals  (Fig.  48,  p.  138)  of  AgCl  are 
deposited  from  the  solution,  and,  under  other  conditions,  crystals  of  the 
composition  2AgCl,  3NH3.  The  bromide,  which  is  less  readily  soluble, 


630  .INORGANIC   CHEMISTRY 

gives  the  same  complex  ion.  The  iodide  is  hardly  soluble  at  all.  Am- 
monio-argention  Ag(NH8)2%  in  solutions  of  concentrations  such  as  are 
commonly  used  (.1 N  to  N),  gives  about  the  same  concentration  of  argen- 
tion  Ag*  as  does  the  bromide,  and  much  more  than  the  highly  insoluble 
iodide  (cf.  p.  544).  Hence  the  latter  is  almost  insoluble  in  ammonium 
hydroxide,  and  can  be  precipitated  in  ammoniacal  solution.  All  three 
of  the  insoluble  halides  dissolve  in  solutions  of  potassium  cyanide  and 
of  sodium  thiosulphate,  as  do  also  all  the  other  insoluble  silver  salts. 
Usually  an  equivalent  amount  of  the  cyanide  or  thiosulphate  suffices, 
but  for  solution  of  the  sulphide  an  excess  is  required.  With  the 
cyanide,  double  decomposition  gives  first  the  insoluble  silver  cyanide 
(AgCN)  which  then  dissolves,  forming  the  soluble  potassium  argenti- 
cyanide  K.Ag(CN)2.  The  thiosulphate  gives  a  solution  from  which 
crystals  of  a  complex  salt  2NaAgS208,  Na2S208  are  obtained.  The  com- 
plex anion  in  the  solution  appears  to  be  Ag(S203)2"'.  Since  the  iodide 
dissolves  in  the  thiosulphate  with  considerable  difficulty,  we  should 
infer  that  the  complex  thiosulphate  anion  gives  about  the  same  concen- 
tration of  argention  as  does  the  iodide.  An  independent  method  of 
measuring  the  concentrations  of  argention  in  all  the  solutions,  places  the 
compounds  in  the  order  of  diminishing  ability  to  give  argention  thus, 
AgCl,  Ag(NH,y,  AgBr,  Ag(S2O3)2'",  Agl,  Ag(CN)/,  Ag2S,  and  con- 
firms  the  above  inferences  (see  Concentration  cells).  The  more  active 
metals,  like  zinc  and  copper,  displace  silver  from  all  solutions,  whether 
the  solutions  contain  simple  or  complex  salts. 

Oxides  of  Silver.  —  When  sodium  or  potassium  hydroxide  is 
added  to  a  solution  of  a  salt  of  silver,  a  pale-brown  precipitate  is 
obtained,  which,  after  being  freed  from  water,  is  found  to  be  Ag2O.  We 
should  expect  to  obtain  the  hydroxide  ( AgOH)  in  this  fashion,  but  it 
appears  to  be  unstable.  The  aqueous  solution  of  argentic  oxide,  how- 
ever, is  distinctly  alkaline,  and  presumably  therefore  does  contain  the 
hydroxide  :  2AgOH  <±  Ag20  -f  H2O.  Silver  oxide  is  formed  by  boil- 
ing silver  chloride  with  caustic  potash.  Since  the  oxide  is  much  more 
soluble  than  the  chloride  (p.  544),  we  should  expect  the  reverse  of  the 
above  action  to  be  the  normal  one.  Here,  however,  the  excess  of  potas- 
sium hydroxide  (hydroxidion)  represses  the  ionization  of  the  silver  hy- 
droxide and  reverses  the  relations  in  regard  to  solubility  (cf.  p.  585). 

Argentic  oxide  melts  and  gives  off  its  oxygen  at  250-270°.  It  is  an 
active  basic  oxide,  and  all  the  salts  of  silver  are  derived  from  it,  the  two 
other  oxides  having  no  corresponding  salts.  When  moist,  it  absorbs 


COPPER,    SILVER,    GOLD  631 

carbon  dioxide  from  the  air.  Its  solutions  are  said  to  show  concentra- 
tions of  hydroxidion  much  smaller,  it  is  true,  than  equimolar  solutions 
of  the  active  bases,  but  considerably  greater  than  similar  solutions  of 
ammonium  hydroxide  (p.  331).  The  oxide  dissolves  easily  in  ammo- 
nium hydroxide,  and  the  ammonio-argentic  hydroxide  Ag(NH8)2.OH 
which  is  formed  is  as  active  a  base  as  is  potassium  hydroxide.  The 
solution,  when  allowed  to  evaporate,  deposits  black  crystals  of  an 
explosive  substance  whose  composition  has  not  been  determined. 
This  is  "  fulminating  silver  "  (not  to  be  confused  with  fulminate  of 
silver  Ag.ONC). 

Silver  peroxide  Ag2O2  (cf.  p.  308)  is  formed  by  the  action  of  ozone 
on  silver.  In  the  electrolysis  of  silver  nitrate  it  is  deposited  in 
shining  black  crystals  on  the  positive  electrode.  There  is  also  a 
suboxide  Ag40. 

Silver  Nitrate.  —  This  salt  is  obtained  by  treating  silver  with 
aqueous  nitric  acid : 

3Ag  +  4HN03  ->  3AgN03  +  NO  +  2H2O. 

From  the  solution,  colorless  rhombic  crystals  (Fig.  7,  p.  13)  isomorphous 
with  those  of  potassium  nitrate  (Fig.  98,  p.  557)  are  deposited.  These 
melt  at  218°.  In  the  form  of  thin  sticks  made  by  casting  (lunar*  caus- 
tic), the  substance  is  used  in  medicine,  partly  because  it  combines  with 
albumins  to  form  insoluble  compounds.  When  commercial  silver,  con- 
taining copper,  is  used  to  make  silver  nitrate,  the  solution  is  evapo- 
rated to  dryness  and  heated  at  250°  until  the  nitrate  of  copper  has  all 
been  decomposed.  At  this  temperature  the  silver  salt  is  unaffected, 
and  when  cool  can  be  separated  from  the  insoluble  cupric  oxide  by  ex- 
traction with  water. 

The  aqueous  solution  is  neutral.  The  pure  salt  is  not  affected  by 
light,  but  when  deposited  on  cloth,  on  the  skin  of  the  fingers,  or  on  the 
mouth  of  the  reagent  bottle,  it  is  converted  into  the  chloride,  and  from 
this,  in  turn,  silver  is  liberated.  For  this  reason  it  is  an  ingredient  in 
marking-inks.  The  dry  compound  combines  with  ammonia,  giving 
AgN08.3NH8.  In  its  aqueous  solution  ammonium  hydroxide  pro- 
duces, first  a  faint  precipitation  of  the  oxide,  and  then  the  soluble 
complex  salt  Ag(NH8)2.N08. 

Other  Salts  of  Silver.  —  Silver  carbonate,  the  neutral  salt  Ag2C08, 
and  not  a  basic  carbonate  is  precipitated  from  solutions  of  salts  of 
*  (Lat.)  luna  (the  moon),  the  alchemical  name  for  silver. 


632  INORGANIC  CHEMISTRY 

silver  by  soluble  carbonates.  It  is  slightly  yellow  in  color.  With  water 
it  gives  a  faint  alkaline  reaction,  and,  like  calcium  carbonate,  is  soluble 
in  excess  of  carbonic  acid  (p.  594).  When  heated,  the  carbonate  decom- 
poses, leaving  metallic  silver.  Other  compounds  of  silver,  for  example, 
the  chloride,  when  heated  in  a  crucible  with  sodium  carbonate  give  this 
salt  by  double  decomposition,  and  hence  are  finally  reduced  to  a  button 
of  metallic  silver.  The  sulphate  is  made  by  the  action  of  concen- 
trated sulphuric  acid  on  the  metal.  It  is  not  very  soluble  in  water,  and 
crystallizes  in  rhombic  prisms  isomorphous  with  anhydrous  sodium 
sulphate.  When  it  is  mixed  with  a  solution  of  aluminium  sulphate 
(q.v.\  octahedral  crystals  of  silver-alum  AggSO^  ALj(S04)3,  24H2O  are 
obtained.  Silver  sulphide  is  precipitated  by  hydrogen  sulphide  from 
solutions  of  all  silver  compounds,  whether  free  acids  are  present  or 
not,  and  irrespective  of  the  form  in  which  the  silver  is  combined. 
Excess  of  potassium  cyanide,  however,  prevents  its  precipitation  from 
the  argenticyanide.  The  sulphide  is  formed  by  the  action  of  metallic 
silver  on  alkaline  hydrosulphides,  and  this  interaction  forms  the  basis 
of  the  «  hepar  "  test  for  sulphur  (p.  391).  Silver  orthophosphate  Ag8PO4 
(yellow),  arsenate  Ag8As04  (brown),  and  chromate  Ag2CrO4  (crimson), 
are  produced  by  precipitation,  and  their  distinctive  colors  enable  us  to 
use  silver  nitrate  in  analysis  as  a  reagent  for  identifying  the  acid 
radicals. 

Electroplating.  —  The  process  is  similar  to  the  electro-deposition 
of  copper  (p.  617).  The  article  to  be  plated  is  cleaned  with  extreme 
care  and  attached  to  the  negative  wire.  A  plate  of  silver  forms  the 
positive  electrode,  and  since  simple  salts  of  silver  do  not  give  coherent 
deposits,  the  bath  is  a  solution  of  potassium  argenticyanide.  The 
kalion  (K*)  migrates  to  the  negative  wire,  and  since  potassium  requires 
a  much  greater  E.M.F.  for  its  liberation  than  does  silver,  silver  is 
there  deposited  from  the  trace  of  argention  given  by  the  complex  silver 
ions  in  the  neighborhood: 


The  potassium  cyanide  remains  in  solution.  At  the  positive  electrode 
silver  goes  into  solution  in  equivalent  amount  giving  argention,  and 
the  above  equations  are  reversed. 

Mirrors  are  silvered  through  the  reduction  of  silver  nitrate  by 
organic  compounds  such  as  potassium-sodium  tartrate  (Rochelle  salt), 


COPPER,    SILVER,   GOLD  633 

glycerine,  formaldehyde  (formol),  or  sugar.  On  a  small  scale,  dilute 
silver  nitrate  is  mixed  with  ammonium  hydroxide  until  the  solution  is 
clear,  and  then  a  little  caustic  potash,  a  few  more  drops  of  ammonia, 
and  finally  a  very  little  glycerine,  are  added.  A  watch-glass  floated 
on  this  mixture  quickly  acquires  a  deposit  of  silver. 

Photography.  —  Bromo-gelatine  dry  plates  are  made  by  preparing 
an  emulsion  of  gelatine  to  which  silver  nitrate  and  a  slight  excess  of 
ammonium  bromide  have  been  added.  After  the  emulsion  has  been 
kept  warm  until  the  precipitate  of  silver  bromide  has  coagulated  into 
small  granules  ("ripening"),  it  is  allowed  to  solidify.  It  is  then  cut 
up,  and  the  ammonium  nitrate  is  washed  out  with  water.  After  drying 
and  remelting,  the  emulsion  is  finally  applied  to  plates  of  glass.  The 
excess  of  ammonium  bromide  and  the  ripening  both  increase  the 
subsequent  sensitiveness  of  the  plates. 

After  exposure,  often  for  only  a  fraction  of  a  second,  there  is  no 
visible  alteration  in  the  film.  The  image  is  developed.  Chemically, 
this  consists  in  reducing  the  silver  bromide  to  metallic  silver  by  means 
of  reducing .  agents.  While  the  whole  of  the  halide  upon  the  plate  is 
reducible,  if  the  reducing  agent  is  kept  upon  it  for  a  sufficient  length 
of  time,  the  parts  reached  by  the  light  are  affected  first,  and  with  a 
speed  proportional  to  the  intensity  of  the  illumination  undergone  by 
each  part.  The  reducing  agent  is  poured  off  when  sunicient "  contrast " 
between  the  parts  variously  illuminated  has  been  attained.  The  un- 
reduced silver  bromide  is  then  dissolved  out  with  sodium  thiosulphate 
("hyposulphite  of  soda"  or  "hypo"),  and  the  silver  image  is  thus 
saved  from  obliteration  by  the  silver  that  would  be  deposited  if 
the  plate  were  to  be  brought  into  the  light  without  this  treatment 
(fixing).  The  result  is  a  "negative,"  as  the  parts  brightest  in  the 
object  are  now  opaque,  and  the  darkest  parts  of  the  object  are  trans- 
parent. 

According  to  one  view,  the  exposure  reduces  certain  portions  of  the 
bromide  to  a  sub-bromide,  perhaps  Ag2Br,  which  is  more  easily  reduced 
than  silver  bromide,  and  is  consequently  first  attacked  by  the  developer. 
The  first  particles  of  free  silver  then  interact  with  neighboring  mole- 
cules of  AgBr,  giving  more  Ag2Br.  Thus  the  reduction  proceeds  exten- 
sively wherever  Ag2Br  is  found  and  in  proportion  to  its  amount.  The 
gelatine  is  the  sensitizing  substance,  and  promotes  the  dissociation  of 
the  silver  bromide  (2AgBr  <=±  Ag2Br  -f-  Br),  which  is  a  reversible  action, 
by  combining  with  the  bromine.  Potassium  bromide,  when  added  to 


634  INORGANIC  CHEMISTRY 

the  developer,  restrains  the  development,  probably  by  rendering  the 
silver  bromide  less  soluble  (cf.  p.  583). 

The  simplest  developer  is  potassium-ferrous  oxalate  K2.Fe(C204)2, 
a  solution  of  which  may  be  made  by  mixing  ferrous  sulphate  and 
potassium  oxalate.  For  the  sake  of  simplicity  we  may  regard  the 
action  as  a  reduction  by  means  of  ferrous  oxalate,  which  itself  is 
oxidized  to  ferric  oxalate  (Fe2(C204)3): 

3FeC204  +  3Ag2Br  ->  Fe2(C2O4)3  +  FeBr3  +  6Ag. 

In  brief,  we  have  3Fe"  becoming  2Fem  +  Fem,  and  this  amount  of 
bivalent  iron  therefore  takes  up  3Br,  liberating  the  silver  with  which 
it  was  combined.  Other  developers  commonly  employed  are  alkaline 
solutions  of  the  sodium  salts  of  hydroquinone  and  pyrogallic  acid. 

In  printing,  the  light  and  dark  are  again  reversed,  the  denser  parts 
of  the  negative  protecting  the  compounds  on  the  paper  below  it  from 
action,  and  leaving  them  white.  Either  "bromide"  papers,  which  require 
only  brief  exposure  and  are  developed  like  the  plate,  are  used,  or  silver 
chloride  is  the  sensitive  substance,  and  prolonged  exposure  to  light  is 
allowed  to  liberate  the  proper  amount  of  silver.  The  operation  of 
fixing  is  performed  as  before.  In  toning,  a  solution  of  sodium  chlor- 
aurate  is  employed.  A  portion  of  the  silver  dissolves,  displacing  gold 
(p.  362),  which  is  deposited  in  its  place : 

NaAuCl4  +  3Ag  -»  NaCl  +  3AgCl  +  Au. 

The  thin  film  of  gold  gives  a  richer  color  to  the  print.  In  platinum 
toning,  potassium  chloroplatinite  K2PtCl4  is  similarly  used. 

Many  other  actions  are  utilized  in  photography.  Thus,  ferric 
oxalate  is  reduced  by  light  to  ferrous  oxalate :  Fe2(C2O4)8  — >  2FeC204 
4-  2C02.  When  paper  coated  with  a  solution  of  the  former,  or  a 
mixture  of  ferric  chloride  and  ammonium  oxalate,  is  used  for  printing, 
the  pale-yellow  ferric  salt  loses  its  color  where  it  has  been  turned  into 
the  ferrous  salt.  If  the  paper  is  then  dipped  in  a  solution  of  ferri- 
cyanide  of  potassium  K3.Fe(CN)6  the  ferrous  salt  precipitates  the 
insoluble  and  deep-blue  ferrous  ferricyanide  Fe3[Fe(CN)6]2,  while  the 
unchanged  ferric  salt  simply  gives  a  soluble  brown  substance,  which 
can  be  washed  out.  For  regular  blue  prints,  ammonium-ferric  citrate  is 
employed  instead  of  the  oxalate.  If  the  above  paper,  after  printing, 
is  dipped  in  potassium  chloroplatinite  (or  has  been  coated  with  this 
salt  at  the  same  time  that  it  received  the  ferric  oxalate),  and  is  then 


COPPER,    SILVER,    GOLD  635 

dipped  in  potassium  oxalate  solution,  the  latter  dissolves  the  insoluble 
ferrous  oxalate,  and  the  potassium-ferrous  oxalate  reduces  the  platinum 
compound,  giving  a  platinum  print : 

6FeC204  +  3K2PtCl4  ->  2Fe2(C204)8  +  2FeCl8  +  3Pt  +  6KC1. 

We  have  already  seen  (p.  484)  that  light  of  short  wave-length  — 
blue  and  violet  —  has  the  greatest  effect  upon  silver  halides.  The 
time,  in  seconds,  required  for  equal  effects  is  approximately :  violet 
15,  blue  29,  green  37,  yellow  330,  red  600.  Hence  objects  showing 
to  the  eye  a  variety  of  colors  are  entirely  misrepresented,  as  regards 
the  relative  brightness  of  their  parts,  by  photography.  Now  the  im- 
portant fact,  in  this  connection  is,  that  only  that  part  of  the  light  which 
is  absorbed  in  traversing  the  film,  and  not  that  which  is  scattered  or 
transmitted,  can  be  used  for  chemical  change.  Hence,  dipping  plates 
in  solutions  of  substances  capable  of  absorbing  yellow  and  red  radia- 
tions causes  them  to  absorb  more  of  the  energy  of  these  photographi- 
cally weakest  radiations,  and  to  give  greater  chemical  action  in  response 
to  them.  This  partially  restores  the  balance.  Such  plates  are  called 
orthochromatic,  and  are  made  with  substances  like  eOsin  or  cyanine. 

Analytical  Reactions  of  Silver  Compounds.  —  The  ion  of  salts 
of  silver,  argention  Ag*,  is  colorless.  Many  of  its  compounds  are  in- 
soluble, the  precipitation  of  the  chloride,  which  is  insoluble  in  dilute 
acids,  being  used  as  a  test.  Mercurous  chloride  and  lead  chloride  are 
also  white  and  insoluble,  but  silver  chloride  dissolves  in  ammonium 
hydroxide,  mercurous  chloride  (<?.v.)  turns  black,  and  lead  chloride, 
which  is  also  soluble  in  hot  water,  is  not  altered  in  color.  With  excess 
of  ammonium  hydroxide,  silver  salts  give  the  complex  cation 
Ag(NH8)2"  and,  from  solutions  containing  silver  in  this  form,  only  the 
iodide  and  sulphide  can  be  precipitated.  Sodium  thiosulphate  and 
potassium  cyanide  dissolve  all  silver  salts,  giving  salts  of  complex  acids 
with  silver  in  the  anion  (p.  629).  Zinc  displaces  silver  from  all  forms 
of  combination. 

GOLD. 

Chemical  Relations  of  the  Element.  —  This  element  forms 
two  very  incomplete  series  of  compounds  corresponding  respectively 
to  aurous  and  auric  oxides,  Au2O  and  Ai^Og.  The  former  is  a 
feebly  basic  oxide,  the  latter  mainly  acid-forming.  No  simple  salts 


636  INORGANIC   CHEMISTRY 

with  oxygen  acids  are  stable.  All  the  compounds  of  gold  are  easily 
decomposed  by  heat  with  liberation  of  the  metal.  All  other  common 
metals  displace  gold  from  solutions  of  its  compounds  (p.  362).  Mild 
reducing  agents  likewise  liberate  gold.  The  element  enters  into  many 
complex  anions  (p.  536). 

Occurrence  and  Metallurgy.  —  Gold  is  found  chiefly  in  the  free 
condition  disseminated  in  veins  of  quartz,  or  mixed  with  alluvial  sand. 
Small  quantities  are  found  also  in  sulphide  ores  of  iron  and  copper. 
Telluride  of  gold  (sylvanite),  in  which  silver  takes  the  place  of  part  of 
the  gold  [Au,Ag]Te2*,  is  found  in  Colorado.  This  mineral  when 
heated  loses  its  tellurium,  and  gold,  alloyed  with  silver,  remains. 

From  the  alluvial  deposits,  gold  is  usually  separated  by  washing 
in  a  cradle,  as  in  the  Klondyke.  Quartz  veins,  which  in  the  Transvaal 
Colony  reach  a  thickness  of  a  meter  and  carry  an  average  of  18  g.  of 
gold  per  ton,  are  mined,  and  the  material  is  pulverized  with  stamping 
machinery.  About  55  per  cent  of  the  gold  is  then  separated  by  allowing 
the  powdered  rock  to  be  carried  by  a  stream  of  water  over  copper  plates 
amalgamated  with  mercury.  The  gold  dissolves  in  the  latter,  and  is 
secured  by  removal  and  distillation  of  the  amalgam.  The  finer  particles 
contained  in  the  sludge  which  runs  off  ("  tailings  "),  are  extracted  by  add- 
ing a  dilute  solution  of  potassium  cyanide  (Mac  Arthur-Forest  process) 
and  exposing  the  mixture  to  the  air.  Oxidation  and  simultaneous  inter- 
action with  the  cyanide  give  potassium  aurocyanide.  Hydrogen  perox- 
ide, which  is  formed  in  many  oxidations  by  free  oxygen,  is  produced 
also: 


2Au  +  4KCN  +  2H2O  +  02  -»  2KAu(CN)2  +  2KOH  +  H202, 
2Au  +  4KCN  +  H202  ->  2KAu(CN)2  +  2KOH. 

From  this  solution  the  gold  is  isolated,  either  by  electrolysis,  in 
which  a  plate  of  lead  forms  the  cathode  (and  is  subsequently  cupelled. 
Siemens-Halske  process),  or  in  the  form  of  a  purple  powder  by  pre- 
cipitation with  zinc. 

Auriferous  pyrites  is  roasted,  and  then  treated  with  chlorine  gas. 
The  chloride  of  gold  which  is  formed  is  dissolved  out  with  water.  From 

*  Amongst  minerals,  mixed  crystals  of  isomorphous  salts  are  so  commonly  found 
that  formulae  like  the  above  are  constantly  used  by  mineralogists.  [Au,Ag]Te, 
indicates  a  mixture  in  varying  proportions  of  the  isomorphous  tellurides  AuTe2  and 
AgTea. 


COPPER,    SILVER,    GOLD  637 

the  solution,  the  gold  is  precipitated  with  ferrous  sulphate  or  oxalic 

acid  : 

2AuCl8  +  6FeS04  -»  2Fea(S04)8  +  2FeCl8  +  2Au, 
2AuCls  +  3H2C204->  6HC1  +  6C02  +  2Au. 

In  the  former  case  a  purple  powder,  and  in  the  latter,  if  the  solution  is 
heated,  a  spongy  mass  (the  form  used  by  dentists),  is  obtained. 

The  gold  separated  from  ores  in  the  above  ways  contains  silver, 
copper,  lead,  and  other  metals,  and  various  methods  of  refining,  elec- 
trolytic and  otherwise,  are  used.  In  one  of  these  the  gold  is  melted, 
and  a  stream  of  chlorine  is  passed  through  it.  The  metals,  excepting 
gold,  are  converted  into  chlorides.  The  chloride  of  silver  rises  as  a 
liquid  to  the  surface,  while  chlorides  of  arsenic  and  antimony  are  vola- 
tilized. A  layer  of  melted  borax  prevents  loss  of  silver  chloride  by 
volatilization.  The  silver  chloride,  when  it  has  solidified,  is  placed 
between  wrought-iron  plates  and  reduced  by  the  action  of  dilute  sul- 
phuric acid  upon  the  latter. 

The  world's  production  of  gold  during  the  first  half  of  the  nine- 
teenth century  averaged  27  tons  annually.  In  1897  it  was-  363  tons, 
and  in  1899, 472.6  tons.  In  the  former  year  North  America,  including 
Canada,  produced  28.5  per  cent  of  the  whole,  the  Transvaal  Colony 
23.2  per  cent,  and  Australia  21.2  per  cent. 


Properties  of  the  Metal»  —  Gold  is  yellow  in  color,  and  is  the 
most  malleable  and  ductile  of  all  the  metals.  It  melts  at  1064°.  To 
give  it  greater  hardness  it  is  alloyed  with  copper,  the  proportion  of  gold 
being  defined  in  "  carats."  Pure  gold  is  "  24-carat."  British  sovereigns 
are  22-carat  and  contain  -fa  of  copper.  American,  French,  and  German 
coins  are  21.6-carat,  or  90  per  cent  gold.  Silver  takes  the  place  of 
copper  in  Australian  sovereigns. 

Gold  is  not  affected  by  free  oxygen  or  by  hydrogen  sulphide.  It 
loes  not  displace  hydrogen  from  dilute  acids,  nor  does  it  interact  with 
nitric  or  sulphuric  acids  or  any  oxygen  acids  except  selenic  acid.  It 
combines,  however,  with  free  chlorine,  and  it  therefore  interacts  with  a 
mixture  of  nitric  and  hydrochloric  acids  (aqua  regia),  which  gives 
off  this  gas  (p.  448).  Chlorauric  acid  H.AuCl4(  =  HC^AuClg)  is 
formed,  and  the  action  is  assisted  by  the  fact  that  the  gold  ions  are 
taken  into  the  little-dissociated  anion  AuCl/.  Gold  is  the  least  active 
of  the  familiar  metals. 


638  INORGANIC  CHEMISTRY 

Compounds  with  the  Halogens.  —  Chlorauric  acid,  formed  as 
above,  is  deposited  in  yellow,  deliquescent  crystals  of  H.AuCl4, 4H20. 
The  yellow  sodium  chloraurate  NaAuCl4, 2H20,  obtained  by  neutraliza- 
tion of  the  acid,  is  used  in  photography  (p.  634).  The  acid  gives  up 
hydrogen  chloride  when  heated  very  gently,  leaving  the  red,  crystalline 
auric  chloride  AuClg.  The  tendency  to  form  complex  compounds  is 
such,  however,  that  when  dissolved  in  water  free  from  hydrochloric 
acid,  this  salt  gives  H2. AuClgO.  Bed  crystals  of  H2AuCl30, 2H20  are 
deposited  by  the  solution.  When  auric  chloride  is  heated  to  180° 
aurous  chloride  AuCl  and  chlorine  are  formed.  This  salt  is  a  white 
powder.  It  is  insoluble  in  water,  but  in  boiling  water  is  converted 
quickly  into  auric  chloride  and  free  gold  :  3AuCl  — >  AuCl3  +  2Au. 
When  potassium  iodide  is  added  to  a  solution  of  chlorauric  acid,  or  to 
sodium  chloraurate,  the  yellow  aurous  iodide  is  precipitated : 

NaAuCl4  +  SKI  ->  NaCl  +  Aul  +  ^  +  3KC1. 

The  action  is  like  that  on  cupric  salts  (p.  621),  and  for  a  similar  reason, 
namely,  that  auric  iodide  is  not  stable. 

Other  Compounds.  —  When  caustic  alkalies  are  added  to  chlor- 
auric acid,  or  to  sodium  chloraurate,  auric  hydroxide  Au(OH)3  is  pre- 
cipitated. This  substance  is  an  acid,  and  interacts  with  excess  of  the 
base,  forming  aurates.  These  are  derived  from  met-auric  acid  (Au(OH)3 
-  H2O  =  HAu02),  as,  for  example,  potassium  aurate  K.Au02, 3H20. 
This  salt  interacts  by  double  decomposition,  giving,  for  instance,  with 
silver  nitrate,  the  insoluble  silver  salt  AgAu02.  Its  solution  is  alkaline 
in  reaction,  a  fact  which  shows  that  auric  acid  is  a  weak  acid  (cf.  p. 
344). 

Auric  oxide  AUgOg  is  a  brown,  and  aurous  oxide  Au^O  is  a  violet 
powder.  With  hydrochloric  acid  the  latter  gives  chlorauric  acid  and 
free  gold. 

On  account  of  its  reducing  action,  hydrogen  sulphide  precipitates 
from  chlorauric  acid  a  dark-brown  mixture  containing  much  aurous 
sulphide  AugS  and  free  sulphur,  as  well  as  some  auric  sulphide  Au^. 
The  sulphides  interact  with  alkali  sulphides,  giving  complex  sulphaurites 
and  sulphaurates,  such  as  K8AuS2(=  3K2S,Au2S)  and  KAuS2(=  K2S, 
AujSg),  which  are  soluble  (p.  537,  and  see  Tin,  Arsenic,  and  Antimony). 

The  aurocyanides,  like  K.Au(CN)2(  =  KCN,  AuCN),  and  the  auri- 
cyanides  like  K.Au(CN)4(=  KCN,  Au(dST)3),  are  formed  by  the  action 


COPPER,   SILVER,    GOLD  639 

of  potassium  cyanide  on  aurous  and  auric  compounds  respectively.  They 
are  colorless  and  soluble.  Their  solutions  are  used  as  baths,  in  con- 
junction with  a  gold  anode,  for  electrogilding. 

Analytical  Reactions  of  Gold.  —  The  metallic  "streak,"  produced 
by  rubbing  the  metal  on  touchstone  (Lydian  stone,  a  black  basalt), 
is  not  easily  removed  by  nitric  acid  of  sp.  gr.  1.36  (57.5  per  cent). 
In  assaying,  the  material  containing  the  gold  is  heated  with  borax  and 
lead  in  a  small  crucible  (cupel)  of  bone-ash.  The  lead  and  copper  are 
oxidized,  and  the  oxides  are  ^ absorbed  by  the  cupel,  leaving  a  drop  of 
molten  alloy  of  gold  and  silver.  The  cold  button  is  flattened  by  ham- 
mering and  rolling,  and  treated  with  nitric  acid  to  remove  the  silver. 
The  gold,  which  remains  unattacked,  is  washed,  fused  again,  and 
weighed.  The  acid  will  not  interact  with  the  silver  and  remove  it 
completely  if  the  quantity  of  gold  exceeds  25  per  cent.  When  the 
proportion  of  gold  is  greater  than  this,  a  suitable  amount  of  pure  silver 
is  fused  with  the  alloy  ("  quartation  "). 

Exercises.  —  1.  How  much  copper  will  be  deposited  per  hour  on 
each  sq.  cm.  of  an  electrode  immersed  in  cupric  sulphate  solution  when 
the  current  density  is  \  ampere  per  sq.  cm.  (p.  323)  ?  How  much 
copper  would  be  obtained  under  the  same  conditions  from  a  cuprous 
salt? 

2.  Write  equations  for  the  interactions  (a)  of  salt  water  and  oxygen 
with  copper  (p.  618),  (b)  of  ferrous  oxide  and  sand  (p.  617),  (c)  of  ver- 
digris, arsenious  acid,  and  acetic  acid  (p.  624). 

3.  Write  the  formulae  of  the  basic  chloride,  nitrate,  carbonates,  and 
sulphate  of  copper  as  if  these  substances  were  composed  of  the  normal 
salt,  the  oxide  and  water  (p.  618). 

4.  What  may  be  the  formula  of  the  compound  of  cupric  hydroxide 
and  sodium  tartrate  (p.  623)  ? 

5.  Can  you  develop  any  relation  between  the  facts  that  solutions 
of  cupric  salts  are  acid  in  reaction  and  that  they  give  basic  carbonates 
by  precipitation  ? 

6.  Formulate  the  action  of  potassium  cyanide  in  dissolving  cupric 
hydroxide    and   cuprous    sulphide,  assuming   that   potassium   cupro- 
cyanide  is  formed. 

7.  How  should  you  set  about  making  cupric  orthophosphate   (in 
solution),  ammonium  cuprocyanide,  and  lead  cuprocyanide  ? 


640  INORGANIC   CHEMISTRY 

8.  Write  the  formulae  of  some  of  the  double  salts  analogous   to 
potassium-cupric  sulphate  (p.  625). 

9.  What  chemical  agents  are  present  in  a  Bunsen  flame  ?     If  borax 
beads  were  made  in  the  oxidizing  flame  with  cupric  chloride,  cuprous 
bromide,   and  cupric  sulphate,    severally,   what   actions   would  take 
place  ? 

10.  If  the  solubility  ratio  of  silver  in  zinc  and  in  lead  were  1000 : 
1,  and  2  per  cent  of  zinc  were  used,  what  proportion  of  the  total  silver 
would  be  secured  by  Parke's  method? 

11.  Which  is  more  stable,  silver  sulphate  or  cupric  sulphate,  silver 
nitrate  or  cupric  nitrate  ?     To  what  salts  are  the  silver  compounds  in 
this  respect  more  closely  allied  ? 

12.  Write  the  equations  for  the  interaction  of  (a)  silver  and  con- 
centrated   sulphuric   acid,  (b)  silver   chloride   and  sodium  carbonate 
when  heated  strongly,  (c)  sodium  thiosulphate  and  silver  bromide,  (d) 
potassium  ferricyanide  and  ferrous  oxalate. 

13.  What  reagents  should  you  use  to  precipitate  the  phosphate, 
arsenate,  and  chr ornate  of  silver  ? 

14.  Write  the'equations  for  the  interactions  of.  (a)  gold  and  selenic 
acid,  in  which  selenious  acid  is  formed,  (b)  potassium  hydroxide  and 
auric  hydroxide,  (c)  potassium  cyanide  and  sodium  chloraurate. 

15.  In  what  respects  are  the  elements  of  this  family  distinctly 
metallic,  and  in  what  respects  are  they  allied  to  the  non-metals  (p.  533)  ? 

16.  Collect  all  the  evidence  tending  to  show  that  the  cuprous  com- 
pounds are  more  stable  than  the  cupric. 

17.  Describe  in  terms  of  the  categories  used  by  the  phase  rule  the 
systems  (a)  cupric  nitrate  and  water  at  24.5°  and  (&)  silver  iodide  at 
146°. 

18.  Make  a  classified  list  of  the  methods  by  which  cupric  com- 
pounds are  transformed  into  cuprous,  and  vice  versa. 


CHAPTER  XXXVII 

GLUCINUM,    MAGNESIUM,      ZINC,    CADMIUM,     MERCURY. 

THE  RECOGNITION  OF    CATIONS  IN  QUALITATIVE 

ANALYSIS 

The  Chemical  Relations  of  the  Family.  —  The  remaining  ele- 
ments of  the  third  column  of  the  periodic  table,  namely,  glucinum  or 
beryllium  (Gl,  or  Be,  at.  wt.  9.1),  magnesium  (Mg,  at.  wt.  24.36),  zinc  (Zn, 
at.  wt.  65.4),  cadmium  (Cd,  at.  wt.  112.4),  and  mercury  (Hg,  at.  wt.  200.0), 
although  all  bivalent,  do  not  form  a  coherent  family.  Glucinum  and 
magnesium  resemble  zinc  and  cadmium,  and  differ  from  the  calcium 
family,  in  that  the  sulphates  are  soluble,  the  hydroxides  easily  lose 
water  leaving  the  oxides,  the  chlorides  are  comparatively  volatile,  and 
the  metals  are  not  rapidly  rusted  in  the  air  and  do  not  easily  displace 
hydrogen  from  water.  They  resemble  the  calcium  family,  and  differ 
from  zinc  and  cadmium,  in  that  the  sulphides  are  hydrolyzed  by  water, 
the  oxides  are  not  reduced  by  heating  with  carbon,  complex  cations  are 
not  formed  with  ammonia,  and  the  metals  do  not  enter  into  complex 
anions.  But  glucinum  differs  from  magnesium  and  resembles  zinc  in 
that  its  hydroxide  is  acidic  as  well  as  basic.  This  is  not  unnatural, 
since  in  the^  periodic  system  it  lies  between  lithium,  a  metal,  and 
boron,  a  non-metal.  Mercury  is  the  only  member  of  the  group  that  forms 
two  series  of  compounds.  These  are  derived  (p.  278)  from  the  oxides 
HgO  and  Hg20.  Mercury  approaches  the  noble  metals  in  the  ease 
with  which  its  oxide  is  decomposed  by  heating,  and  in  the  position  of 
the  free  element  in  the  electromotive  series. 

The  vapor  densities  of  zinc,  cadmium,  and  mercury  show  the 
vapors  of  these  three  metals  to  be  monatomic. 

The  compounds  of  the  metals  of  this  family  give  no  color  to  the 
borax  bead. 

GLUCINUM. 

Chemical  Relations  of  the  Element.  —  Glucinum  (or  beryllium) 
is  bivalent  in  all  its  compounds.  Its  oxide  and  hydroxide  are  basic, 
and  are  also  feebly  acidic  towards  active  bases  (see  Zinc  hydroxide). 

641 


642  INORGANIC   CHEMISTRY 

On  account  of  this  fact  and  the  extreme  ease  with  which  its  carbonate 
gives  up  carbon  dioxide,  in  both  of  which  respects  it  resembles  alu- 
minium, it  was  first  thought  to  be  trivalent.  This  made  its  atomic 
weight  13.6,  the  amount  combining  with  one  chemical  unit  of  chlorine 
being  4.55.  In  the  periodic  system,  however,  there  was  a  space  for  a 
bivalent  element  with  the  atomic  weight  9.1  (=  2  x  4.55)  between 
lithium  and  boron,  and  none  for  a  trivalent  element.  Later  (1884) 
Nilson  and  Pettersson  determined  the  vapor  density  of  the  chloride  and 
of  certain  organic  compounds  of  the  element,  and  found  only  9.1  parts 
of  glucinum  in  the  molar  weights  of  the  compounds.  The  element 
derives  its  name  from  the  sweet  taste  of  its  salts  (Gk.  yXv/cv?,  sweet). 


The  Metal  and  its  Compounds.  —  Glucinum  occurs  in  beryl,  a 
metasilicate  of  glucinum  and  aluminium  AL^Gl^SiOg),..  Specimens  of 
beryl  tinted  green  by  the  presence  of  a  little  silicate  of  chromium  are 
known  as  emeralds.  The  metal  may  be  obtained  by  electrolysis  of 
the  easily  fusible  double  fluoride  G1F2,  2KF.  In  powdered  form  it 
burns  when  heated  in  the  air.  It  displaces  hydrogen  from  cold,  dilute 
acids,  and  also,  when  heated,  from  caustic  potash  : 

Gl  +  2KOH  ->  K2G102  +  H2. 

The  oxide  interacts  with  acids  and  with  strong  bases.     The  salts  give 
no  color  to  the  Bunsen  flame. 

MAGNESIUM. 

Chemical  Relations  of  the  Element.  —  Magnesium  is  bivalent 
in  all  its  compounds.  The  oxide  and  hydroxide  are  basic  exclusively. 
The  element  does  not  enter  into  complex  cations  or  anions. 

Occurrence.  —  Magnesium  carbonate  occurs  alone  as  magnesite, 
and  in  a  double  salt  with  calcium  carbonate  MgCO8,  CaCO8  as  dolomite. 
The  sulphate  and  chloride  are  found  as  hydrates  and  as  constituents 
of  double  salts  (see  below)  in  the  Stassfurt  deposits.  Silicates  are 
also  common.  Olivine  is  the  orthosilicate  Mg2Si04.  Serpentine  is  a 
hydrated  disilicate,  [Mg,Fe]3,Si207,  2H20,  as  is  also  meerschaum,  and 
asbestos  is  an  anhydrous  silicate.  The  element  derives  its  name  from 
Magnesia,  a  town  in  Asia  Minor. 

The  Metal.  —  Magnesium  is  manufactured  by  electrolysis  of  de- 
hydrated and  fused  carnallite  MgCl2,KCl,  6H20.  The  iron  crucible 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY         643 

in  which  the  material  is  melted  forms  the  cathode,  and  a  rod  of  carbon 
the  anode.  The  metal  is  silver-white,  and  when  heated  can  be  pressed 
into  wire  and  rolled  into  ribbon.  Commercial  specimens  of  the  latter 
often  contain  zinc. 

Chemically  the  metal  is  less  active  than  are  the  metals  of  the  alka- 
line earths.  It  slowly  becomes  coated  with  a  layer  of  the  oxide.  It 
displaces  hydrogen  from  boiling  water  and,  of  course,  from  cold,  dilute 
acids.  Magnesium  burns  in  air  with  a  white  light,  rich  in  rays  of 
short  wave-length  such  as  act  upon  photographic  plates  (p.  458).  The 
ash  contains  the  nitride  Mg3lN~2,  as  well  as  the  oxide.  The  presence 
of  the  former  may  be  shown  by  the  evolution  of  ammonia  when  the 
white  powder  is  boiled  with  water  (p.  417).  When  the  metal  is  heated 
with  the  oxides  of  boron,  of  silicon,  and  of  many  of  the  metals,  it  com- 
bines with  the  oxygen  and  liberates  the  other  element. 

Powdered  magnesium  is  used  in  pyrotechny,  and,  with  potassium 
chlorate  (10  :  17),  in  making  flash-light  powder  for  use  in  photography. 

Magnesium  Chloride.  —  This  salt  occurs  in  salt  deposits  as  the 
hexahydrate  MgCLj,  6H20,  a  highly  deliquescent  compound  obtained  also 
by  evaporating  an  aqueous  solution,  and  as  carnallite  MgCLj,  KC1,  6H20. 
The  latter  is  an  important  source  of  potassium  chloride  (p.  551),  and 
almost  all  the  magnesium  chloride  combined  with  it  is  thrown  away. 
When  the  hexahydrate  is  heated,  a  part  of  the  chloride  is  hydrolyzed, 
some  magnesium  oxide  remaining,  and  some  hydrogen  chloride  being 
given  off.  Sea-water  cannot  be  used  in  ships'  boilers  because  of  the 
hydrochloric  acid  liberated  by  the  magnesium  chloride  which  the  water 
contains.  The  salt  forms  a  double  chloride  with  ammonium  chloride 
MgCLj,  NH4C1,  6H20  which  is  isomorphous  with  carnallite,  and  this  salt 
can  be  dehydrated  without  hydrolysis  of  the  chloride.  Afterwards 
the  ammonium  chloride  can  be  volatilized  (p.  421).  To  utilize  natural 
magnesium  chloride,  the  manufacture  of  chlorine  from  it,  by  passing 
air  and  steam  over  the  salt  at  a  high  temperature,  has  been  attempted  : 

2H20  +  O2  ->  4MgO  +  4HC1  +  2CL,. 


The  Oxide  and  Hydroxide.  —  Magnesium  oxide  is  made  by 
heating  the  'carbonate,  and  is  known  as  "calcined  magnesia."  It  is  a 
white,  highly  infusible  powder,  and  is  used  for  lining  electric  furnaces 
and  making  crucibles.  It  combines  slowly  with  water  to  form  the 
hydroxide. 


644  INORGANIC   CHEMISTRY 

The  hydroxide  is  found  in  nature  as  brucite.  It  is  also  precipi- 
tated from  solutions  of  magnesium  salts  by  alkalies.  It  is  very 
slightly  soluble  in  water,  much  less  so  than  calcium  hydroxide,  but 
more  so  than  are  the  hydroxides  of  zinc  and  the  other  heavy  metals. 
The  solution  has  a  barely  perceptible  alkaline  reaction. 

Magnesium  hydroxide  is  not  precipitated  by  ammonium  hydroxide 
when  ammonium  salts  are  present  also.  The  ammonium  salts,  being 
highly  ionized  and  giving  a  high  concentration  of  ammonion  NH4*,  re- 
press the  ionization  of  the  feebly  ionized  ammonium  hydroxide,  and  so 
reduce  the  concentration  of  hydroxidion  which  it  furnishes.  With  the 
ordinary  concentration  of  Mg",  therefore,  the  amount  of  hydroxidion 
existing  in  presence  of  excess  of  a  salt  of  ammonium  is  too  small  to 
bring  the  solubility  product  [Mg**]  x  [OH']2  up  to  the  value  required 
for  precipitation.  Conversely,  magnesium  hydroxide  interacts  with 
solutions  of  ammonium  salts  and  passes  into  solution  : 

Mg(OH)2  (solid)  <=>  Mg(OH)2  (diss'd)  «=»  Mg"+  20H' 


2NH4C1  <±  2C1'  '    **       H* 


In  presence  of  excess  of  ammonium  chloride,  the  OH'  combines  with 
NH*4  to  form  molecular  ammonium  hydroxide,  and  the  equilibria  in  the 
upper  line  are  displaced  forwards  to  generate  a  further  supply  of  the 
former.  With  sufficiently  great  concentration  of  the  ammonium  chloride, 
all  the  magnesium  hydroxide  may  thus  dissolve  ;  with  only  a  small  excess 
a  condition  of  equilibrium  with  solid  magnesium  hydroxide  is  reached. 
The  whole  case  is  analogous  to  the  interaction  of  acids  with  insoluble 
salts  (p.  598).  Magnesium  oxide  also  dissolves  in  salts  of  ammonium. 
It  gives  first  the  hydroxide  by  interaction  with  the  water. 

Magnesium  Carbonate.  —  The  normal  carbonate  is  found  in  nature. 
Only  hydrated  basic  carbonates  are  formed  by  precipitation,  and  their 
composition  varies  with  the  conditions.  The  carbonate  manufactured  in 
large  amounts  and  sold  as  magnesia  alba  is  approximately  Mg4(OH)2 
(C08)8.  3H20.  The  carbonates  are  not  precipitated  in  the  presence  of 
ammonium  salts,  and  interact  with  such  salts  in  the  same  way  as  does 
the  hydroxide. 

Magnesium  Sulphate.  —  The  common  heptahydrate  MgS04,  7H20 
crystallizes  from  cold  water  in  rhombic  prisms,  and  is  called  Epsom  salts. 
At  0°  a  dodecahydrate  appears.  The  heptahydrate  is  efflorescent,  and 
loses  its  water  by  stages  and  with  decreasing  aqueous  tension.  The  mono- 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY        645 

hydrate,  found  in  the  salt  layers  as  kieserite  MgS04,H20,  has  a  very 
low  aqueous  tension,  and  is  not  rapidly  dehydrated  except  above  200°. 
(The  hepta-  and  inonohydrates  present  a  striking  case  of  difference  in 
solubility  in  two  forms  of  one  salt,  the  former  giving  at  15°  a  solution 
containing  33.8  g.  of  the  sulphate  in  100  g.  of  water,  while  the  latter 
is  almost  insoluble.  Magnesium  sulphate  is  used  in  the  manufacture 
of  sodium  and  potassium  sulphates,  and  is  employed  also  for  "  loading  " 
cotton  goods,  and  as  a  purgative. 

Magnesium  Sulphide.  —  The  sulphide  may  be  formed  by  heating 
the  metal  with  sulphur.  It  is  insoluble  in  water,  but  is  decomposed 
and  gives,  finally,  hydrogen  sulphide  and  magnesium  hydroxide  : 

2MgS  4-  2H20  ±*Mg(SH)2  +  Mg(OH)2, 
Mg(SH)2  +  2H,0  ^  Mg(OH)2 1  +  2H2S. 

The  hydrolysis  is  more  complete  than  in  the  case  of  calcium  sulphide, 
and  eliminates  all  the  hydrogen  sulphide,  because  magnesium  hydrox- 
ide is  much  more  insoluble  than  calcium  hydroxide,  and  so  there  is 
little  reverse  interaction  tending  to  reproduce  the  soluble  hydro- 
sulphide  Mg(SH)2. 

Phosphates  of  Magnesium.  —  The  only  phosphate  of  importance 
is  ammonium-magnesium  orthoph'osphate  NH4MgPO4, 6H20,  which 
appears  as  a  crystalline  precipitate  when  sodium  phosphate  and 
ammonium  hydroxide  are  mixed  with  a  solution  of  a  magnesium  salt. 
This  compound  is  insoluble  in  water  containing  ammonium  hydroxide, 
and  is  used  in  quantitative  analysis  for  estimating  both  magnesium 
and  phosphoric  acid.  Before  being  weighed  the  precipitate  is  ignited, 
and  is  thus  converted  into  the  anhydrous  pyrophosphate  of  magnesium 
Mg2P207.  The  salt  NH4MgAs04,6H20  has  similar  properties,  and  is 
used  for  estimating  arsenic  acid. 

Analytical  Reactions  of  Magnesium  Compounds.  —  The  mag- 
nesium ion  is  colorless  and  bivalent.  It  does  not  enter  into  complex 
ions.  Soluble  carbonates  precipitate  basic  carbonates  of  magnesium, 
but  not  when  ammonium  salts  are  present.  The  latter  limitation  dis- 
tinguishes compounds  of  magnesium  from  those  of  the  calcium  family. 
Potassium  hydroxide  precipitates  the  hydroxide  of  magnesium,  except 
when  salts  of  ammonium  are  present.  The  mixed  phosphate  of  ammo- 
nium and  magnesium,  in  presence  of  ammonium  hydroxide,  is  the 
least  soluble  salt. 


646  INORGANIC   CHEMISTRY 

ZINC. 

Chemical  Relations  of  the  Element.  —  Zinc  is  bivalent  in  all  its 
compounds.  Of  these  there  are  two  sets,  —  the  more  numerous  and 
important  one  in  which  zinc  is  the  positive  radical  (Zn.S04,  Zn.Cl2, 
etc.),  and  a  less  numerous  set,  the  zincates,  in  which  zinc  is  in  the 
negative  radical  (Na2.Zn02,  etc.).  Both  sets  of  salts  are  hydrolyzed  by 
water,  as  the  hydroxide  is  feeble  whether  it  is  considered  as  an  acid  or 
as  a  base.  The  element  also  enters  into  complex  cations  and  anions. 
The  salts  are  all  poisonous. 

Occurrence  and  Extraction  from  the  Ores.  —  The  chief  sources 
of  zinc  are  calamine  or  smithsonite  ZnC03,  zinc-blende  (G-er.  blenden, 
to  dazzle)  or  sphalerite  ZnS,  franklinite  Zn(Fe02)2,  and  zincite  ZnO. 
The  red  color  of  the  last  is  due  to  the  presence  of  manganese. 

The  ores  are  first  converted  into  oxide  —  the  carbonate  by  ignition, 
and  the  sulphide  by  roasting.  The  sulphur  dioxide  is  used  to  make 
sulphuric  acid.  A  mixture  of  the  oxide  wrth  coal  is  then  distilled  in 
earthenware  retorts  at  1300-1400°,  the  zinc  condensing  in  earthenware 
receivers,  while  carbon  monoxide  burns  at  a  small  opening : 

2ZnS  +  302  ->  2ZnO  +  2S02, 

ZnO  +  C  -+C0  +  Zn. 

At  first  zinc  dust,  a  mixture  of  zinc  and  zinc  oxide,  collects  in  the 
receiver,  and  afterwards  liquid  zinc.  The  product,  which  is  cast  in 
blocks,  is  called  spelter.  It  contains  small  amounts  of  lead,  arsenic, 
iron,  and  cadmium,  because  the  sulphides  of  these  metals  are  almost 
invariably  present  in  zinc-blende. 

Properties  and  Uses  of  the  Metal.  —  Zinc  is  a  bluish-white 
crystalline  metal.  When  cold  it  is  brittle,  but  at  120-150°  it  can  be 
rolled  into  sheets  between  heated  rollers  and  then  retains  its  pliability 
when  cold.  At  200-300°  the  metal  becomes  once  more  brittle,  at  433° 
it  melts,  and  at  920°  it  boils.  The  vapor  density  at  1740°  is  2.64,  and 
the  molecular  weight,  therefore,  2.64  x  28.955  (p.  214)  or  76.4. 
The  gas  is  thus  monatomic. 

The  metal  burns  in  air  with  a  bluish  flame,  giving  zinc  oxide. 
When  cold  it  is  not  affected  by  dry  air,  but  in  moist  air  it  is  oxidized, 
and  becomes  covered  with  a  firmly  adhering  layer  of  basic  carbonate 
which  protects  it  from  further  action.  The  metal  displaces  hydrogen 
from  dilute  acids,  but  with  pure  specimens  the  action  almost  ceases  in 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY        647 

consequence  of  the  formation  of  a  layer  of  condensed  hydrogen  on  the 
surface.  Contact  with  a  less  electro-positive  metal,  such  as  lead,  iron, 
copper,  or  platinum,  enables  the  action  to  go  on,  because  the  hydrogen 
is  then  liberated  at  the  surface  of  the  other  metal  (see  Electromotive 
chemistry).  Crude  zinc  contains  lead  and  iron  and  is  therefore  more 
active  than  pure  zinc.  Zinc  also  attacks  boiling  alkalies,  giving  the 
soluble  zincate  (see  below)  :  2KOH  +  Zn  -^K2Zn02  +  H2.  The  action 
on  ammonium  hydroxide  is  slower  and  different  in  nature  : 

Zn  +  2NH4OH  +  2NH8  ->  Zn(NH8)4.(OH)2+  H2, 

a  complex  cation  being  formed. 

Sheet  zinc,  in  consequence  of  its  lightness  (sp.  gr.  7),  is  used  in 
preference  to  lead  (sp.  gr.  11.5)  for  roofs,  gutters,  and  architectural 
ornaments.  Galvanized  iron  is  made  by  dipping  cleaned  sheet  iron  in 
molten  zinc.  The  latter,  being  more  active  (p.  362),  is  rusted  instead 
of  the  iron.  Zinc  is  used  also  in  batteries  and  for  making  alloys 
(p.  619).  It  mixes  in  all  proportions  with  tin,  copper,  and  antimony, 
but  with  lead  (p.  627)  and  with  bismuth  separation  into  two  layers 
occurs,  each  metal  dissolving  only  a  little  of  the  other.  The  two 
different  modes  of  behavior  resemble  those  of  alcohol  and  water 
(p.  136)  and  ether  and  water  (p.  147)  respectively. 

Zinc  Chloride.  —  This  salt  is  usually  manufactured  by  treating 
zinc  with  excess  of  hydrochloric  acid,  evaporating  the  solution  to  dry- 
ness,  and  fusing  the  residue.  When  hydrochloric  acid  is  thus  present, 
the  chloride  ZnCl^  is  obtained.  Evaporation  of  the  pure  aqueous  solu- 
tion, which  is  acid  in  reaction,  results  in  considerable  hydrolysis  and 
formation  of  much  of  the  basic  chloride 


H2O  «=±  HC1  +  Zn(OH)Cl,  (1) 

2Zn(OH)Cl  ->  Z^OCL,  +  H2O.  (2) 


The  salt  is  used  in  solid  form  as  a  caustic  and,  by  injection  of  a  solu- 
tion into  wood  (e.g.,  railway  sleepers),  as  a  poison  to  prevent  the 
growth  of  organisms  which  promote  decay.  In  both  cases  the  salt  com- 
bines with  albumins,  forming  solid  products.  The  aqueous  solution, 
being  acid,  is  employed  also  for  dissolving  the  oxides  from  surfaces 
which  are  to  be  soldered.  The  acid  is  reproduced  by  hydrolysis  as  fast 
as  it  is  used,  and  finally  the  oxychloride  remains  (equation  1  above). 


648  INORGANIC   CHEMISTRY 

Zinc  Oxide  and  Hydroxide  and  the  Zincates.  —  The  oxide  is  ob- 
tained as  a  white  powder  by  burning  zinc  or  by  heating  the  precipitated 
basic  carbonates.  It  turns  yellow  when  heated,  recovering  its  white- 
ness when  cold,  in  the  same  way  that  mercuric  oxide  is  brown  whilst 
hot  and  bright  red  when  cold.  It  is  employed  in  making  a  paint  — 
zinc-white  or  Chinese  white  —  which  is  not  darkened  by  hydrogen 
sulphide.  For  filling  teeth,  dentists  sometimes  use  a  paste  made  by 
mixing  the  oxide  with  a  strong  solution  of  zinc  chloride.  It  quickly 
sets  to  a  hard  mass  of  oxychloride. 

The  hydroxide  of  zinc  appears  as  a  white,  flocculent  solid  when 
alkalies  are  added  to  solutions  of  zinc  salts.  It  interacts  as  a  basic 
hydroxide  with  acids,  giving  salts  of  zinc  : 

Zn(OH)2  -h  H2S04  <±  Zn.S04  +  2H2O. 

It  also  interacts  with  excess  of  the  alkali  employed  to  precipitate  it, 
giving  a  soluble  zincate : 

H2Zn02f  +  2KOH  <=±  K2.ZnO2  +  2H2O. 

Both  actions  are  reversible,  and  the  second  requires  a  considerable 
excess  of  alkali  for  its  completion  :  in  fact,  most  of  the  zinc  hydroxide 
seems  to  be  simply  in  colloidal  solution.  From  a  consideration  of  these 
facts  it  is  evident  that  zinc  hydroxide  when  in  solution  is  ionized  both 
as  an  acid  and  as  a  base : 

2H*  +  ZnO,"4=±  Zn(OH)2  (diss'd)  <=»  Zn-+  20H' 
Zn(OH)2  (solid) 

The  ionization  as  an  acid  is  less  than  that  as  a  base,  but  both  are 
small.  Addition  of  an  acid  like  sulphuric  acid,  however,  furnishes 
hydrion ;  the  hydroxyl  ions  combine  with  this  to  form  water,  and  all 
the  equilibria  are  displaced  to  the  right.  With  a  base,  on  the  other 
hand,  the  hydrion  is  removed  and  the  basic  ionization  simultaneously 
repressed,  so  that  the  equilibria  are  displaced  to  the  left. 

Zinc  hydroxide  interacts  with  ammonium  hydroxide,  giving  a  soluble 
complex  compound  with  ammonia  Zn(NH3)4.(OH)2.  The  case  is  like 
those  of  copper  (p.  623)  and  silver  hydroxides  (p.  631),  and  not  like 
that  of  magnesium  hydroxide  (p.  644). 

Compounds  of  zinc,  when  heated  in  the  Bunsen  flame  with  a  salt 
of  cobalt,  gives  a  zincate  of  cobalt  (Rinmann's  green)  CoZnCX,. 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,   MERCURY         649 

Hydrogen  sulphide  precipitates  zinc  sulphide  from  solutions  of 
zincates  and  from  solutions  containing  ammonia,  so  that  some  zinc 
ions  Zn"  are  present  in  both. 

Carbonate  of  Zinc.  —  The  normal  zinc  carbonate  may  be  precipi- 
tated by  means  of  sodium  bicarbonate  : 


ZnS04  +  2NaHC03  _>  Na,SO4  +  ZnCO3  +  H20  +  CO2. 

The  normal  carbonate  of  sodium,  however,  gives  basic  carbonates, 
which,  as  in  the  case  of  magnesium  (p.  644),  vary  in  composition 
according  to  the  conditions. 

Zinc  Sulphate.  —  This  salt  is  formed  when  zinc-blende  is  roasted. 
It  gives  rhombic  crystals  of  the  hydrate  ZnS04,  THgO.  This,  and  the 
corresponding  compounds  of  magnesium  MgS04,  TtLjO,  of  iron  FeS04, 
7H20,  and  of  several  other  bivalent  metals,  are  all  isomorphous,  and  are 
known  as  vitriols.  The  zinc  salt  is  white  vitriol.  Like  Epsom  salts, 
it  is  dehydrated  by  stages,  the  last  molecule  of  water  being  difficult  to 
remove.  It  is  used  in  cotton-printing  and  as  an  eye-wash  (£  per  cent 
solution). 

The  salt  gives  double  salts  with  potassium  and  ammonium  sulphate,  of  , 
the  form  ZnS04,K2S04,6H20,  which  crystallize  in  the  monosynimetric 
system,  and  are  isomorphous  with  each  other,  and  with  double  salts 
containing  copper  (p.  625),  mercury  (Hg11),  iron  (Fe11),  magnesium,  and 
other  bivalent  elements  in  place  of  the  zinc.  These  compounds,  unlike 
the  complex  cyanides,  are  almost  completely  decomposed  in  dilute 
solution  (cf.  p.  537). 

Zinc  Sulphide.  —  This  compound  is  the  only  familiar  sulphide 
which  is  white.  The  yellow  color  of  zinc-blende  is  caused  by  the 
presence  of  sulphide  of  iron.  Zinc  sulphide  is  more  soluble  in  water 
than  is  sulphide  of  copper,  and  hence  it  interacts  with  excess  of  strong 
acids,  and  passes  into  solution.  It  is  not  soluble  enough,  however,  to 
be  much  affected  by  weak  acids  like  ace  tic  acid..  This  sort  of  behavior 
is  shown  also  by  calcium  oxalate  (p:  598),  and  was  discussed  fully  in  that 
connection.  Zinc  sulphide  is  thus  capable  of  being  precipitated  when 
acetic  acid  is  present,  or  when  hydrogen  sulphide  is  led  into  a  solution 
of  the  acetate  of  zinc  : 

Zn(C2H302)2  +  H2S  <=>  ZnS  J  +  2HC2H802. 


650  INORGANIC  CHEMISTRY 

But  when  an  active  acid  is  present,  or  is  formed  during  the  operation, 
the  sulphide  is  precipitated  incompletely  or  not  at  all,  the  action  being 
highly  reversible  : 

ZnS04  +  H2S  <=»  ZnS  +  H2S04. 

There  are  thus  two  ways  of  obtaining  the  sulphide  by  precipita- 
tion. A  soluble  sulphide  causes  it  to  be  thrown  down  completely 
because  no  acid  is  liberated  in  the  action : 

ZnCl2  +  (NH4)2S  <±  ZnSJ  +  2NH4C1. 

The  other  method  is  to  add  sodium  acetate  to  the  solution  of  the  salt, 
and  then  lead  in  hydrogen  sulphide.  The  acid  which  is  liberated  by  the 
action  upon  the  salt  interacts  with  the  sodium  acetate,  giving  a  neutral 
salt  of  sodium  and  acetic  acid,  and  the  zinc  sulphide  is  not  affected  by 
the  latter.  In  terms  of  the  ionic  hypothesis,  the  hydrion,  liberated  as 
the  hydrogen  sulphide  interacts  with  the  zinc  salt,  combines  with 
acetanion  introduced  by  the  sodium  acetate,  and  gives  the  little-ion- 
ized acetic  acid. 

Analytical  Reactions  of  Zinc  Salts.  —  Zinc  sulphide  is  precipi- 
tated by  the  addition  of  ammonium  sulphide  to  solutions  of  zinc  salts 
and  of  zincates.  Sodium  hydroxide  gives  the  insoluble  hydroxide, 
which,  however,  interacts  with  excess  of  the  alkali,  giving  the  soluble 
zincate  of  sodium.  Compounds  of  zinc,  when  heated  on  charcoal  with 
cobalt  nitrate,  give  Rinmanu's  green  (p.  648). 

CADMIUM. 

Chemical  Relations  of  the  Element.  —  This  element  is  biva- 
lent in  all  its  compounds.  Its  oxide  and  Itydroxide  are  basic  exclu- 
sively, and  the  salts  are  not  hydrolyzed  by  water.  It  enters  into  com- 
plex compounds  having  the  ions  Cd(NH3)4**,  Cd(CN)4",  and  CdI4". 

The  Metal.  —  Aside  from  the  rare  mineral  greenockite  CdS,  cad- 
mium is  found  only  in  small  amounts,  as  carbonate  and  sulphide,  in  the 
corresponding  ores  of  zinc.  During  the  reduction,  being  more  volatile 
than  zinc,  it  distils  over  first. 

In  color  the  metal  resembles  tin,  and  is  much  more  malleable  and 
ductile  than  zinc.  It  melts  at  320°  and  boils  at  770°. 

It  displaces  hydrogen  from  dilute  acids  but  is  itself  displaced  from 
solutions  of  its  compounds  by  zinc,  since  it  is  less  electro-positive. 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY        651 

Compounds  of  Cadmium.  —  The  chloride  crystallizes  as  a  dihy- 
drate  CdCl2, 2H20,  which  is  efflorescent  and  is  not  hydrolyzed  during 
dehydration  or  in  solution.  Zinc  chloride  (p.  647)  is  deliquescent  and 
is  easily  hydrolyzed.  The  halides  are  less  ionized  than  are  the  corre- 
sponding compounds  of  most  other  metals.  The  iodide,  in  particular, 
seems  to  exist  in  solution  as  Cd.CdI4,  and  the  complex  anion  gives 
little  ionic  cadmium.  On  account  of  this  fact  even  the  sulphide  can- 
not be  precipitated  completely  from  a  solution  of  the  iodide.  Con- 
versely, hydriodic  acid  dissolves  the  sulphide  to  a  much  greater  extent 
than  do  other  acids  (see  below). 

The  hydroxide  is  made  by  precipitation,  and  interacts,  as  a  basic 
hydroxide,  with  acids,  but  not  at  all  with  bases.  It  dissolves  in  ammo- 
nium hydroxide,  however,  forming  Cd(NH3)4,(OH)2.  The  oxide  is  a 
brown  powder,  obtained  by  heating  the  hydroxide,  carbonate,  or  nitrate, 
or  by  burning  the  metal. 

The  sulphate  crystallizes  from  solution  as  3CdS04,8H20,  and  is 
not  isomorphous  with  the  sulphates  of  zinc  and  magnesium.  Soluble 
carbonates  throw  down  the  normal  carbonate  of  cadmium  CdC03,  and 
not  a  basic  carbonate. 

Hydrogen  sulphide  precipitates  the  yellow  sulphide  CdS  even  from 
acid  solutions  of  the  salts.  The  substance  is  used  as  a  pigment. 
Since  zinc  sulphide  is  not  formed  under  these  conditions  (p.  650), 
the  first  part  of  the  distillate  from  the  reduction  of  the  ore  can  be 
dissolved  in  hydrochloric  acid  and  the  cadmium  precipitated  as 
sulphide,  while  the  zinc  remains  in  solution.  The  sulphide  of 
cadmium,  however,  is  less  insoluble  in  water  than  are  the  sulphides 
of  copper  and  mercury,  and  therefore  cannot  be  precipitated  from  a 
strongly  acid  solution.  We  have  thus  a  distinct  gradation  in  the 
solubility  in  water  of  various  insoluble  sulphides.  The  order  of 
increasing  insolubility,  and  consequent  difference  in  behavior  towards 
acids,  is : 

Barium,  strontium,  cai-  -\ 

cium,  and  magnesium  V  attacked  by  water,  and  all  acids. 

sulphides :  ) 

Ferrous  sulphide  :  attacked  even  by  acetic  acid,  but  not  by  water. 

Zinc  sulphide  :  attacked  by  dilute  active  acids,  but  not  by  acetic  acid. 

Cadmium  sulphide  :         I  attacked  by  concentrated  active  acids,  but  not  by  dilute 

/      acids. 

Cupric  sulphide  :  \  attacked  by  fairly  concentrated,  actively  oxidizing  acids, 

|      like  nitric  acid,  but  not  much  by  ordinary  active  acids. 
Mercuric  sulphide  :  hardly  attacked  at  all,  even  by  hot  nitric  acid. 


652  INORGANIC   CHEMISTRY 

Analytical  Reactions  of  Cadmium  Compounds.  —  The  cad- 
mium ion  Cd**  is  bivalent  and  colorless.  The  yellow  cadmium  sul- 
phide is  precipitated  by  hydrogen  sulphide,  even  from  acid  solutions 
of  the  salts.  It  is  also  precipitated  from  solutions  containing  the 
complex  cation  Cd(NH3)4"  and  the  complex  anion  Cd(CN)4",  as,  for 
example,  from  a  solution  made  by  adding  excess  of  potassium  cyanide 
to  cadmium  chloride  solution  (K2.Cd(CN)4).  The  latter  property 
enables  cadmium  to  be  separated  from  copper  (p.  624).  The  white 
hydroxide  is  thrown  down  by  sodium  hydroxide,  and  is  not  soluble  in 
excess  of  this  reagent.  It  is  not  formed  from  solutions  containing  the 
Cd(NH8)4**  and  Cd(CN)4"  ions,  and  dissolves  in  ammonium  hydroxide. 
These  and  other  precipitations  are  not  complete  when  cadmium  iodide 
Cd.CdI4  is  used. 

MERCURY. 

Chemical  Relations  of  the  Element.  —  Like  copper,  this  element 
enters  into  two  series  of  compounds,  the  mercurous  and  the  mercuric, 
in  which  it  is  univalent  and  bivalent  respectively.  The  mercurous 
halides,  like  the  cuprous  halides  (and  the  argentic  halides),  are  insoluble 
in  water  and  are  decomposed  by  light.  There  are,  however,  mercurous 
as  well  as  mercuric  salts  of  oxygen  acids.  Both  of  the  oxides,  Hg20  and 
HgO,  are  basic  exclusively,  but  in  a  feeble  degree.  The  hydroxides,  like 
silver  hydroxide,  are  not  stable,  and  lose  water,  giving  the  oxides.  The 
salts  of  both  sets  are  markedly  hydrolyzed  by  water,  and  basic  salts  are 
therefore  common.  No  carbonate  is  known.  Mercury  enters  into  the 
anions  of  a  number  of  complex  salts,  such  as  HgCl4",  HgI4",  Hg(CN)4", 
etc.  It  does  not  give  complex  cations  with  ammonia  resembling  those 
of  cadmium,  copper,  and  silver  (Cd(NH3)4",  etc.),  from  which  ammonia 
is  removed  by  heating,  but  instead  forms  a  class  of  mercur-ammonium 
compounds  like  HgnNH2Cl,  all  of  which  are  insoluble.  The  mercuric 
halides  and  cyanide  show  many  peculiarities  due  to  their  being  very 
little  ionized.  Salts  as  a  class  are  highly  ionize^  bodies,  and  those  of 
mercury  and,  to  a  less  degree,  those  of  cadmium  are  the  only  con- 
spicuous exceptions. 

The  mercury  salts  of  volatile  acids,  like  the  corresponding  salts  of 
ammonium  (p.  421),  can  all  be  volatilized  completely.  Mercury  vapor 
and  all  mercury  compounds  are  poisonous,  the  soluble  ones  more 
markedly  so  than  the  insoluble  ones. 

The  mercurous  salts,  as  a  rule,  are  formed  when  excess  of  mercury 
is  employed,  and  mercuric  salts  when  excess  of  the  oxidizing  acid  or 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY         653 

other  substance  is  present.  Reducing  agents  turn  mercuric  into  mer- 
curous  salts,  and  oxidizing  agents  do  just  the  reverse. 

As  in  the  case  of  the  cuprous  compounds,  it  is  a  question  whether 
simple  or  multiple  formulae,  HgCl  or  HgjGLj,  etc.,  should  be  employed 
for  mercurous  salts.  Pending  the  discovery  of  some  basis  for  a  decision, 
the  simple  formulae  are  used  here. 

Occurrence  and  Isolation  of  the  Metal.  —  Mercury  occurs  to 
some  extent  native  and  to  a  larger  extent  as  red,  crystalline  cinna- 
bar, mercuric  sulphide  HgS.  The  chief  mines  are  in  Spain,  California, 
and  Austria. 

The  liberation  of  the  metal  is  easy,  because  the  sulphide  is  decom- 
posed at  a  high  temperature,  and  the  sulphur  forms  sulphur  dioxide. 
The  mercury  does  not  unite  with  oxygen,  for  the  oxide  decomposes 
(p.  12)  at  400-600° : 

HgS  +  02  ->  Hg  +  S02. 

In  some  places  the  ore  is  spread  on  perforated  brick  shelves  in  a  verti- 
cal furnace,  and  the  gases  pass  through  tortuous  flues  in  which  the 
vapor  of  the  metal  condenses.  The  product  is  filtered  through  chamois- 
skin.  For  separation  from  metallic  impurities,  like  zinc,  arsenic,  and 
tin,  which  are  dissolved,  it  must  be  distilled.  In  the  laboratory,  where 
mercury  finds  many  applications,  it  becomes  impure  with  use,  and  then 
adheres  to  glass,  and  does  not  run  freely  in  spherical  droplets.  For 
purification  it  is  placed  along  with  a  little  diluted  nitric  acid  in  a 
separatory  funnel  (Fig.  56,  p.  147),  and  kept  in  continual  agitation  by 
means  of  a  current  of  air  drawn  or  blown  through  the  mass.  By  this 
treatment  foreign  metals,  such  as  sodium  or  zinc,  nearly  all  of  which 
are  much  more  active  than  mercury  (cf.  p.  362),  are  converted  into 
nitrates.  Pure,  dry  mercury  can  be  drawn  off,  when  needed,  at  the 
bottom.  If  a  high  degree  of  purity  is  required,  the  product  must  be 
distilled  in  vacua. 

Physical  Properties.  —  Mercury  or  quicksilver  (NL.  hydrargy- 
rum^ from  Gk.  v8«/>,  water,  and  apyvpos,  silver)  is  a  silver-white  liquid. 
At  —  39.5°  it  freezes,  and  at  357°  it  boils.  At  ordinary  temperatures  it 
has  a  measurable  vapor  tension,  at  15°  0.0008  mm.  and  at  99°  0.26  mm. 
The  vapor  is  colorless,  does  not  conduct  electricity,  and  is  monatomic. 
A  gold-leaf  suspended  over  mercury  becomes  amalgamated,  since  the 
solution  of  gold  in  mercury  has  a  vapor  tension  smaller  than  that  of 
pure  mercury  (p.  161). 


654  INORGANIC  CHEMISTRY 

On  account  of  its  high  specific  gravity  (13.6,  at  0°)  and  low  vapor 
tension,  the  metal  is  employed  for  filling  barometers  and  manometers. 
Its  uniform  expansion  favors  its  use  in  thermometers.  The  tendency 
to  form  amalgams,  which  it  exhibits  towards  all  the  familiar  metals  with 
the  exception  of  iron  and  platinum  (the  latter,  however,  is  "wet"  by 
it),  is  taken  advantage  of  in  various  ways.  Sodium  amalgam  (p.  569), 
which  is  solid  when  the  sodium  exceeds  2  per  cent,  behaves  like  free 
sodium,  but  with  moderated  activity.  A  layer  of  mercury  on  the  zinc 
plates  of  batteries  reduces  the  action  of  the  acid  on  the  zinc,  while  the 
cells  are  not  in  use.  Mixtures  of  mercury  with  powdered  cadmium 
and  a  small  proportion  of  copper,  quickly  form  solid  amalgams,  and  are 
used  by  dentists.  The  employment  of  mercury  in  the  extraction  of  gold 
by  washing  has  already  been  mentioned  (p.  636).  Mirrors  backed  with 
a  tin-mercury  amalgam  have  been  almost  completely  displaced  by  silvered 
mirrors  (p.  632). 

Chemical  Properties. — When  kept  at  a  temperature  near  to  its 
boiling-point,  mercury  combines  slowly  with  oxygen.  Its  inactivity 
towards  oxygen  when  cold  places  it  next  to  the  noble  metals.  On 
account  of  its  general  inactivity,  it  is  used  in  the  laboratory  for  confin- 
ing gases.  It  does  interact  with  hydrogen  sulphide  and  hydrogen 
iodide,  however  (cf.  Silver,  p.  628).  Mercury  does  not  displace  hydro- 
gen from  dilute  acids  (p.  362),  but  with  oxidizing  acids  like  nitric  acid 
and  hot  concentrated  sulphuric  acid,  the  nitrates  and  sulphates  are 
formed.  With  excess  of  mercury,  the  mercurous  salts,  and  with  ex- 
cess of  the  hot  acid,  the  mercuric  salts,  are  produced.  When  triturated, 
for  example  with  milk-sugar,  mercury  is  divided  into  minute  droplets 
with  relatively  large  surface.  In  this  form  it  is  used  in  medicine 
("blue  pills"),  and  shows  an  activity  which  is  entirely  wanting  in 
larger  masses. 

Mercurous  Chloride.  —  This  salt  (calomel)  is  obtained  as  a  white 
powder  by  precipitation  from  solutions  of  mercurous  salts.  It  is  man- 
ufactured by  subliming  mercuric  chloride  with  mercury  : 

HgCl2  +  Hg->2HgCl, 

or  more  usually  by  subliming  a  mixture  of  mercuric  sulphate,  made  as 
described  above,  with  mercury  and  common  salt.  It  is  deposited  on  the 
cool  part  of  the  vessel  as  a  fibrous  crystalline  mass,  or,  when  the  vapor  is 
led  into  a  large  chamber,  as  a  fine  powder.  It  is  slowly  affected  by  light 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY         655 

just  as  silver  chloride  is.  Here,  however,  the  chlorine  which  is  released 
combines  with  another  molecule  of  the  salt  to  form  mercuric  chloride. 
Since  the  vapor  pressure  of  calomel  reaches  760  mm.  before  the  tem- 
perature has  risen  to  the  melting-point,  the  compound  sublimes  at 
atmospheric  pressure  without  melting.  Its  vapor  density  corresponds 
to  the  formula  HgCl,  but  the  vapor  can  be  shown  to  contain  more  or 
less  Hg  -f-  HgCLj.  Since  this  change  does  not  alter  the  average 
molecular  weight,  and  the  action  is  reversed  when  the  temperature 
falls  (cf.  Ammonium  chloride,  p.  421),  it  has  not  been  found  possible 
to  determine  whether  little  or  much  of  the  calomel  is  thus  affected. 
The  substance  is  used  in  medicine  on  account  of  its  tendency  to  stim- 
ulate all  organs  producing  secretions. 

Mercuric  Chloride.  —  By  direct  union  with  chlorine  the  mercuric 
salt,  corrosive  sublimate  HgCLj,  is  formed.  It  is  usually  manufac- 
tured by  subliming  mercuric  sulphate  with  common  salt,  and  crystal- 
lizes in  white,  rhombic  prisms.  It  melts  at  265°  and  boils  at  307°.  At 
20°  one  hundred  parts  of  water  dissolve  only  7.4  parts  of  the  salt,  and 
at  100°,  54  parts.  It  is  more  soluble  in  alcohol  and  in  ether.  The 
aqueous  solution  is  slightly  acid  in  reaction.  The  salt  is  easily  reduced 
to  mercurous  chloride.  When  excess  of  stannous  chloride  is  added  to 
the  solution,  the  white  precipitate  of  calomel,  first  formed,  passes  into 
a  heavy  gray  precipitate  of  finely  divided  mercury  : 


*  SnCl4  +  2HgCl, 
2HgCl  4-  SnCL,  ->  SnCl4  +  2Hg. 

The  halides  of  mercury  are  very  little  ionized  in  solution,  the 
bromide  and  iodide  even  less  so  than  the  chloride.  Hence  these  salts 
are  little  affected  by  sulphuric  acid  or  nitric  acid.  For  example,  the  chlo- 
rine and  nitrosyl  chloride  which  hydrochloric  acid  forms  with  the  nitric 
acid  are  not  observed  when  mercuric  chloride  is  added  to  this  acid.  On 
this  account,  too,  the  hydrolysis  of  the  chloride  is  much  less  than  that 
of  the  nitrate.  There  is  a  tendency  also  to  the  formation  of  complex 
salts,  so  that  the  addition  of  sodium  chloride  increases  the  solubility 
in  water  and  renders  the  solution  neutral,  NaCl,  HgCl2  or  NaHgCl8  being 
formed.  The  complex  salts,  like  K.HgCl3,H20,  H2.HgCl4,7H20,  and 
NH^HgC^HO,  are  easily  made  by  crystallization  from  solution.  The 
anions  are  relatively  highly  ionized,  however,  and  the  behavior  is  inter- 
mediate between  that  of  complex  salts  and  double  salts  (p.  537). 


656  INORGANIC   CHEMISTRY 

Corrosive  sublimate,  when  taken  internally,  is  extremely  poisonous. 
A  very  dilute  solution  is  used  in  surgery  to  destroy  lower  organisms 
and  thus  prevent  infection  of  wounds.  The  pharmaceutical  tabloids 
of  mercuric  chloride  contain  sodium  chloride,  because,  although  the 
latter  diminishes  the  activity  of  the  compound,  it  also  does  away  with 
the  formation  of  insoluble  chlorides  and  hastens  solution.  Mercuric 
chloride  acts  also  as  a  preservative  of  zoological  materials,  forming  in- 
soluble compounds  with  albumins,  and  preventing  their  decay.  For 
the  same  reason,  albumin  (white  of  an  egg)  is  given  as  an  antidote  in 
cases  of  sublimate  poisoning. 

The  Iodides  of  Mercury.  —  Mercurous  iodide  is  formed  by  rub- 
bing iodine  with  excess  of  mercury.  It  also  appears  as  a  greenish- 
yellow  precipitate  when  potassium  iodide  is  added  to  a  solution  of  a 
mercurous  salt.  The  compound  decomposes  spontaneously  into  mer- 
cury and  mercuric  iodide.  The  decomposition  is  much  hastened  by  the 
use  of  excess  of  potassium  iodide,  which  combines  with  and  removes  the 
mercuric  iodide  (see  below)  : 

2HgI  <±  Hg  +  HgI2. 

Mercuric  iodide  is  obtained  by  direct  union  of  mercury  with  excess 
of  iodine,  or  by  addition  of  potassium  iodide  to  a  solution  of  a  mercuric 
salt.  It  is  a  scarlet  powder,  insoluble  in  water,  but  soluble  in  alcohol 
and  ether.  It  interacts  with  excess  of  potassium  iodide,  forming  the 
soluble,  colorless  potassium  mercuri-iodide  K2.HgI4  with  which  many 
precipitants  fail  to  give  mercury  compounds.  When  heated  above 
116.5°  it  turns  slowly  into  a  yellow  modification,  and  at  223°  this  new 
form  melts.  On  being  cooled,  it  freezes  first  in  the  tetragonal  yellow 
form,  and  below  116.5°,  especially  if  touched  with  a  glass  rod,  it 
changes  into  the  red,  monoclinic  variety  with  evolution  of  heat.  Sul- 
phur (p.  368)  and  ammonium  nitrate  (p.  565)  show  similar  transition 
points.  When  the  vapor  of  the  compound  is  cooled,  it  first  forms  thin 
scales  of  the  yellow  form,  which  is  the  unstable  one  at  low  tempera- 
tures, and  these  turn  red  when  touched.  Similarly,  precipitation  gives 
first  the  yellow  variety,  which  presently  becomes  red  (cf.  Transforma- 
tion by  steps,  p.  453). 

The  Oxides.  —  When  bases  are  added  to  solutions  of  mercurous 
salts,  a  brownish-black  powder  is  thrown  down,  which,  when  dried,  is 
found  to  be  mercurous  oxide  Hg20.  The  hydroxide  is  doubtless 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY         657 

formed  transitorily  and  then  loses  water  (cf.  Silver  oxide,  p.  630). 
Under  the  influence  of  light  or  gentle  heat  (100°),  this  compound  re- 
solves itself  into  mercuric  oxide  and  mercury. 

Mercuric  oxide  HgO  is  formed  as  a  red,  crystalline  powder,  when 
mercury  is  heated  in  air  near  to  357°,  but  is  usually  made  by  decompos- 
ing the  nitrate.  Commercial  specimens,  incompletely  decomposed,  thus 
frequently  give  nitrogen  tetroxide  when  heated.  It  is  formed  also  as  a 
yellow  powder  by  adding  bases  to  solutions  of  mercuric  salts.  It  is 
contended  by  some  chemists  that  the  difference  in  activity  between  the 
red  and  yellow  forms  is  due  solely  to  the  finer  state  of  division  of  the 
latter,  and  by  others  it  is  maintained  that  the  substance  is  dimorphous 
(p.  369)  and  that  two  distinct  varieties  exist. 

The  Nitrates.  —  The  mercurous  salt  is  formed  by  the  action  of 
cold,  diluted  nitric  acid  upon  excess  of  mercury.  It  forms  monoclinic 
crystals  of  a  hydrate  HgN03,H2O.  It  is  hydrolyzed,  slowly  by  cold, 
and  rapidly  by  warm  water,  giving  a  basic  nitrate  : 

2HgN03  +  H20  +±  HN08  +  Hg2(OH)N03  \  . 

On  this  account  a  clear  solution  can  be  made  only  when  some  nitric 
acid  is  added.  Free  mercury  is  also  kept  in  the  solution  to  reduce 
mercuric  nitrate,  which  is  formed  by  atmospheric  oxidation  : 

or     Hg"+  Hg  ->  2Hg\ 


Mercuric  nitrate  is  produced  by  using  excess  of  warm,  concen- 
trated nitric  acid.  It  forms  rhombic  tables  of  Hg(N03)2,8H20.  The 
aqueous  solution  is  strongly  acid,  and  deposits  a  yellowish,  crystalline, 
basic  nitrate  Hg3(OH)20(N03)2.  The  action  •  is  reversed  by  adding 
nitric  acid. 

Sulphides  of  Mercury.  —  Mercurous  sulphide  Hg2S  is  formed 
by  precipitation  from  mercurous  salts,  but  is  stable  only  below  —  10°. 
Above  this  temperature  it  decomposes  into  mercury  and  mercuric 
sulphide. 

Crystallized  mercuric  sulphide  occurs  as  cinnabar,  and  is  red. 
When  formed  by  precipitation  with  hydrogen  sulphide,  or  by  rubbing 
together  mercury  and  sulphur,  it  is  black  and  amorphous.  By  subli- 
mation, in  the  course  of  which  it  dissociates,  the  black  form  gives  the 
red,  crystalline  one.  When  allowed  to  stand  under  a  solution  of  so- 
dium sulphide,  the  black  form  is  slowly  transformed  into  the  red.  This 


658  INORGANIC   CHEMISTRY 

shows  that,  as  we  should  expect,  the  red  form  is  the  more  stable, 
possesses  less  energy,  and  is  less  soluble  at  ordinary  temperatures. 
The  change  is  effected  by  intermediate  formation  of  a  complex  sul- 
phide, the  solution,  when  saturated  toward  the  less  stable  black  sul- 
phide, being  supersaturated  toward  the  more  stable  red  one.  A  white, 
crystalline  sodium  mercuri-sulphide  NagHgSg,  8H20  can,  in  fact,  be  ob- 
tained from  the  solution. 

The  black  and  the  red  varieties  do  not  interact  with  concentrated 
acids,  and  are  even  unaffected  by  boiling  nitric  acid,  which  oxidizes 
most  sulphides  readily.  They  are,  therefore,  less  soluble  in  water 
than  cupric  sulphide  (pp.  600,  651),  and  much  less  so  than  sulphide 
of  cadmium.  They  are  attacked,  however,  by  aqua  regia. 

The  red  form  of  the  sulphide  is  used  in  making  paint  (vermilion). 
The  color  is  more  permanent  than  that  of  red  lead  (Pb304),  because  re- 
ducing gases  (e.g.  S02),  acids  (e.g.  H2S04),  and  hydrogen  sulphide, 
which  are  present  in  the  air,  do  not  affect  it.  It  is  not  stable,  how- 
ever, when  applied  to  metals,  since  iron,  zinc,  etc.,  all  displace  mer- 
cury from  combination,  and  in  these  cases,  therefore,  red  lead  is  pre- 
ferred. 

Mercuric  Cyanide.  — This  salt  is  made  by  treating  precipitated 
mercuric  oxide  with  hydrocyanic  acid,  and  is  obtained  in  square-pris- 
matic crystals.  When  heated  it  gives  off  cyanogen  :  Hg(CN)2  — > 
Hg  -f-  C2N2,  and  is  a  convenient  source  of  this  gas  (p.  507).  The  com- 
pound is  soluble  irvalcohol,  ether,  and  water.  In  solution  in  water 
it  is  so  little  ionized  that  the  freezing-point  of  the  solution  is  normal 
(p.  292),  and  many  reagents  fail  to  show  the  presence  of  either  ion. 
Thus,  with  silver  nitrate  no  silver  cyanide  is  precipitated,  and  with 
a  base  no  mercuric  oxide.  With  potassium  cyanide  it  forms  a  complex 
cyanide  K2.Hg(CN)4.  Hydrogen  sulphide  throws  down  the  sulphide 
from  both  the  simple  and  the  complex  cyanides. 

The  Fulminate  and  Thiocyanate-  —  Mercuric  fulminate 
Hg(ONC)2  is  obtained  as  a  white  precipitate  when  mercury  is  treated 
with  nitric  acid,  and  alcohol  is  added  to  the  solution.  It  decomposes 
suddenly  when  struck,  and  is  used  in  making  percussion  caps.  The 
thiocyanate  Hg(SCN)2  is  precipitated  when  potassium  thiocyanate 
K(SCN)  is  added  to  a  solution  of  a  mercuric  salt.  When  formed  into 
little  balls  and  burned  in  the  air,  the  substance  leaves  a  curiously 
voluminous  ash  ("Pharaoh's  serpents"). 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY        659 

Mercur-ammonium  Compounds.  —  When  ammonium  hydrox- 
ide is  added  to  a  solution  of  a  mercuric  salt,  a  white  substance,  of  a  type 
which  we  have  not  previously  encountered,  is  thrown  down.  Mercuric 
chloride  gives  HgNH2Cl,  commonly  called  "infusible  white  precipi- 
tate," and  the  nitrate  gives  HgNH2N08 : 

HgCl2  +  2NH3  ->  HgNH2Cl  +  NH4G1. 

The  substance  may  be  regarded  as  being  derived  from  ammonium 
chloride  (p.  421)  by  the  displacement  of  2H  by  Hg : 

/H 

Hg:N-H. 
\C1 

It  would  thus  be  named  mercur-ammonium  chloride. 

When  the  mercuric  chloride  is  added  to  a  boiling  solution  of 
ammonium  chloride  containing  ammonium  hydroxide,  mercur-diam- 
monium  chloride  Hg(NH3Cl)2,  "  fusible  white  precipitate,"  appears. 

The  addition  of  ammonium  hydroxide  to  a  solution  of  potassium 
niercuri-iodide  K2HgI4  gives  rise  to  a  third  type  of  compound, 
dimercur-ammonium  iodide  Hg2NI,H2O,  which  appears  as  a  brown 
precipitate : 

2HgI2  +  4NH3  -»  Hg2NI  +  3NH4L 

A  solution  of  potassium  mercuri-iodide  containing  potassium  hydroxide, 
Nessler's  reagent,  becomes  distinctly  yellow  with  traces  of  ammonia, 
and  brown  with  larger  amounts,  and  is,  therefore,  a  valuable  reagent 
for  detecting  traces  of  this  base. 

When  calomel  is  treated  with  ammonium  hydroxide,  it  turns  into  a 
black,  insoluble  body.  This  appears  to  be  a  mixture  of  free  mercury, 
to  which  it  owes  its  dark  color,  and  "infusible  white  precipitate," 
Hg  +  HgNH2Cl.  To  this  reaction  calomel  owes  its  name  (Gk. 
KaXo/xeA.as,  beautiful  black).  Mercurous  nitrate  gives  a  black,  insoluble 
mixture,  Hg  +  HgNH2NO3.  Calomel,  when  dry,  absorbs  ammonia  gas, 
forming  a  molecular  compound  of  the  common  type  HgCl,  NH3.  This 
substance  loses  the  ammonia  again  when  the  pressure  is  reduced.  The 
other  compounds  described  above,  on  the  other  hand,  do  not  contain 
nitrogen  and  hydrogen  in  the  proportions  necessary  to  form  ammonia 
and  are  stable.  Hence  they  are  necessarily  to  be  regarded  as  belong- 
ing to  a  different  type. 

Analytical  Reactions  of  Mercury  Compounds.  —  The  two  ionic 
forms  of  the  element,  monomercurion  Hg*  and  diniercurion  Hg  ",  are 


660  INORGANIC  CHEMISTRY 

both  colorless,  but  their  chemical  behavior  is  entirely  different.  Both 
give  the  black  sulphide  HgS,  which  is  insoluble  in  acids  and  other 
solvents  of  mercury  salts.  Monomercurion  gives  the  insoluble,  white 
chloride,  the  black  oxide,  and  a  black  mixture  with  ammonium 
hydroxide.  Dimercurion  gives  a  soluble  chloride,  a  yellow,  insoluble 
oxide,  and  a  white  precipitate  with  ammonium  hydroxide.  The  be- 
havior with  stannous  chloride  (p.  655)  is  characteristic  of  mercuric  salts. 
With  potassium  iodide  the  behavior  of  the  two  ions  is  quite  different 
(p.  656).  The  more  active  metals  displace  mercury  from  all  compounds. 
This  displacement  is  employed  as  a  test  for  compounds  of  mercury. 
Copper  is  used  by  preference  as  the  displacing  metal  because  the  mer- 
cury is  easily  seen  on  its  surface. 

Salts  of  mercury  are  volatile.  When  heated  in  a  tube  with  sodium 
carbonate,  they  also  give  a  sublimate  of  metallic  mercury. 

THE  RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS. 

"  Wet-way  "  analysis  consists  in  recognizing  the  various  positive  and 
negative  ions  present  in  a  solution  (p.  343).  In  discussing  hydrogen 
sulphide  (p.  375),  it  was  stated  that  the  sulphides  might  be  divided 
into  three  classes  according  to  their  behavior  towards  water  and  acids. 
Now  these  differences  in  behavior  furnish  us  with  a  basis  for  distin- 
guishing the  cations  present  in  a  solution.  Since  the  properties  of  a 
number  of  sulphides  and  other  compounds  of  the  metals  have  been 
studied  in  recent  chapters,  it  is  possible  to  make  a  more  complete 
statement. 

The  following  plan,  taken  in  conjunction  with  the  statements  in 
the  context,  shows  how  a  single  cation  may  be  identified,  and  how. 
when  several  cations  are  present,  a  separation  preparatory  to  identifi- 
cation may  be  effected.  What  will  be  said  applies  only  to  the  case  of 
a  solution  containing  salts  like  the  chlorides,  nitrates,  or  sulphates  of 
one  or  more  cations,  and  leaves  the  oxalates  (p.  601),  phosphates, 
cyanides,  and  some  other  salts,  out  of  consideration. 

Group  1.  —  It  is  usual  to  add,  first,  hydrochloric  acid,  to  find  out 
whether  cations  giving  insoluble  chlorides  are  present.  Argentic, 
m,ercurous,  and  plumbic  salts  give  the  white  AgCl,  HgCl,  and  PbCl2 
respectively.  (For  the  further  recognition  of  each,  see  p.  635.)  Fil- 
tration eliminates  the  precipitate,  if  there  is  any. 

Group  2.  —  A  free,  active  acid  being  now  present,  hydrogen  sulphide 
is  led  into  the  solution.  The  sulphides  insoluble  in  active  acids, 


RECOGNITION   OF  CATIONS  661 

namely,  HgS,  CuS,  PbS,  Bi2S3,  CdS,  As2S3,  Sb2S3>  SnS,  are  therefore 
thrown  down.  The  first  four  are  black  or  brown,  the  next  two  are 
yellow,  and  the  last  two  are  orange  and  brown  respectively.  A  dark- 
colored  substance  will  naturally  obscure  one  of  lighter  color,  if  more 
than  one  is  present.  If  too  much  acid  is  used,  the  precipitation  of 
several  of  the  sulphides  will  be  incomplete  (p.  651) ;  if  too  little,  zinc 
sulphide  may  come  down  (p.  649).  Filtration  again  eliminates  the 
precipitate. 

This  group  is  easily  subdivided.  Any  or  all  of  the  last  threa  sul- 
phides will  pass  into  solution  when  warmed  with  yellow  ammonium 
sulphide,  for  they  give  soluble  complex  sulphides  similar  to  potassium 
sulphaurate  (p.  638).  The  first  five  sulphides,  or  any  of  them,  will  be 
unaffected.  On  the  other  hand,  these  five  sulphides,  with  the  exception 
of  HgS,  will  interact  with  hot  nitric  acid  (p.  658).  Other  reactions 
described  in  the  context  are  then  used  to  distinguish  between,  or,  if 
there  is  a  mixture,  to  separate,  the  members  of  the  sub-groups. 

Group  3.  —  The  solution  (filtrate)  is  now  neutralized  with  ammonium 
hydroxide,  and  ammonium  sulphide  is  added.  Some  ammonium  chlo- 
ride is  also  used,  to  prevent  the  precipitation  of  magnesium  hydroxide 
(p.  644),  which,  in  any  event,  would  be  incomplete.  The  sulphides 
which  are  insoluble  in  water,  and  are  not  hydrolyzed  by  it,  now  appear. 
They  are  FeS,  CoS,  NiS,  all  black,  MnS,  H2O,  and  ZnS,  which  are  pink 
and  white  respectively.  There  are  precipitated  also  the  hydroxides  of 
chromium  and  of  aluminium,  Cr(OH)3  and  A1(OH)3,  because,  although 
their  sulphides  are  hydrolyzed  by  water,  the  hydroxides  are  formed  by 
the  hydroxidion  in  the  ammonium  sulphide  solution.  They  are  too 
insoluble  to  behave  like  magnesium  hydroxide  (p.  646)  by  dissolving 
in  salts  of  ammonium.  They  also  form  no  complex  metal-ammonia 
ion,  as  does  zinc  (p.  648).  The  sulphides  of  nickel  and  cobalt  re- 
semble the  sulphide  of  zinc  in  being  precipitated  by  hydrogen  sulphide 
when  acetic  acid  is  the  only  acid  present.  The  other  sulphides  inter- 
act even  with  acetic  acid  (p.  651). 

Another  plan  is  to  oxidize  the  iron,  if  present,  and  use  ammonium 
chloride  and  ammonium  hydroxide  instead  of  ammonium  sulphide. 
The  hydroxides  of  the  trivalent  elements,  Fe(OH)8,  Gr(OH)3,  A1(OH)8, 
can  be  precipitated  by  excess  of  ammonium  hydroxide  even  when  salts 
of  ammonium  are  present.  Those  of  the  bivalent  metals,  Mn(OH)2, 
Fe(OH)2,  Zn(OH)2,  M(OH)2,  Co(OH)2,  resemble  magnesium  hydroxide 
(p.  644),  and,  of  these,  the  last  three  resemble  also  zinc  hydroxide 
(p.  648),  and  so  cannot  be  precipitated.  After  filtration,  ammonium 


662  INORGANIC   CHEMISTRY 

sulphide  now  throws  down  the  sulphides  of  the  five  bivalent  metals 
(for  a  third  plan,  see  Chemical  relations  of  aluminium). 

Group  4.  —  After  filtration  from  members  of  the  iron  group,  if  any 
were  present,  ammonium  carbonate  is  added,  and  precipitates  the  re- 
maining metals  whose  carbonates  are  insoluble,  BaCO3,  SrCO3,  CaCO3, 
with  the  exception  of  magnesium  (p.  644). 

By  addition  of  sodium  phosphate  to  a  portion  of  the  filtrate,  mag- 
nesium, if  present,  now  comes  out  in  the  form  NH4MgPO4.  There  re- 
main in  solution  only  salts  of  potassium,  sodium,  and  ammonium. 
Since  only  ammonium  compounds,  and  other  substances  which  can  be 
volatilized  have  been  added,  evaporation  and  ignition  of  the  residue 
leaves  the  salts  of  the  two  metals.  If  no  other  metallic  elements  have 
been  shown  to  be  present,  it  saves  time  to  examine  a  fresh  portion  of  the 
original  material.  Salts  of  ammonium  must  also  be  sought  in  a  fresh 
sample  by  the  usual  test  (p.  421). 

The  following  simple  compounds  are  soluble,  but  are  so  little 
ionized  that  their  solutions  do  not  show  all  the  reactions  of  both  of 
the  ions :  NH/)H,  H2S,  HNC,  H2C08,  HgCl2,  Hg(CN)2,  Fe(CNS)3, 
Fe(C2H8O2)2.  With  a  number  of  others,  for  example  CdI2,  the  actions 
are  incomplete  for  the  same  reason.  Complex  compounds,  as  we  have 
seen,  give  complex  ions,  and  the  latter  are  usually  so  little  resolved 
into  simpler  ions  that  the  latter  cannot  be  discovered  by  all  the  usual 
tests.  Thus :  K.Ag(CN)2  gives  K*  and  Ag(CN)/,  but  very  little  Ag* 
and  CN'  (p.  629) ;  Cu^H^Cl,  gives  much  Cu(NH8)2",  and  Cl'  but 
very  little  Cu**.  The  individual  cases  are  fully  described  in  the 
context. 

Exercises.  —  1.  What  is  the  numerical  value  (a)  of  the  solubility 
product  of  magnesium  hydroxide,  (b)  of  the  concentration  of  hydrox- 
idion  given  by  it  and  by  normal  ammonium  hydroxide  respectively 
(p.  544)  ?  Will  normal  concentration  of  ammonium  chloride  suffice  to 
reduce  the  latter  below  the  former  ? 

2.  Why  should  we,  perhaps,  expect  ammonium  sulphide  solution  to 
precipitate  magnesium  hydroxide,  and  why  does  it  not  do  so  ? 

3.  What  volume  of  air  is  required  to  oxidize  one  formula-weight 
of  zinc  sulphide  to  ZnO  and  S02,  and  what  volume  of  sulphur  dioxide 
is  produced?     Is  the  product  more  or  less  diluted  with  nitrogen  than 
when  pure  sulphur  is  burned,  and  by  how  much  ? 

4.  Make  equations  showing  (a)  the  effect  of  heating  zinc  chloride 
with  cobalt  nitrate  (Co(N08)2)  in  the  Bunsen  flame  (p.  648),  (b)  the 


RECOGNITION  OF  CATIONS  663 

action  of  hydrogen  sulphide  on  sodium  zincate,  (c)  the  actions  of  con- 
centrated nitric  acid  and  of  concentrated  sulphuric  acid  on  mercury. 

5.  What  is  the  distinction  between  a  solid  isomorphous  mixture  of 
two  salts  and  a  double  salt  ? 

6.  What  kind  of  salts  might  take  the  place  of  sodium  acetate  in 
the  precipitation  of  zinc  sulphide  (p.  650)  ?     Give  examples. 

7.  Compare  the  amalgamation  of  a  gold-leaf  by  mercury  vapor  with 
the  phenomenon  of  deliquescence  (p.  162). 

8.  If  the  scheme  for  the  recognition  of  cations  (p.  660)  were  ap- 
plied to  solutions  prepared  from  materials  containing  (a)  calcium  oxa- 
late  and  (b)  potassium  argenticyanide,  at  what  stage  and  how  would 
the  presence  of  each  of  these  substances  affect  the  normal  order  ? 

9.  Why  do  none  of  the  salts  of  the  elements  in  this  family  give 
recognizable  effects  with  the  borax  bead  ? 


CHAPTER   XXXVIII 
ELECTROMOTIVE  CHEMISTRY 

WE  have  seen  that  chemical  changes  which  proceed  spontaneously 
liberate  some  form  of  energy.  As  a  rule,  heat  is  developed  ;  but  with 
special  arrangement  of  the  apparatus  (p.  20)  so  that  it  takes  the 
form  of  a  battery-cell,  an  equivalent  amount  of  electricity  is  obtained 
instead.  To  avoid  suggesting  that  this  energy  comes  from  nothing, 
we  say  that  the  original  system  contained  a  certain  amount  of  free,  or 
available,  chemical  energy  and  that  this  has  been  transformed,  con- 
comitantly  with  the  chemical  change,  into  an  equivalent  amount  of 
heat,  or  of  electrical  energy,  as  the  case  may  be.  Nor  is  the  alterna- 
tive of  producing  electrical  energy  available  only  when  the  action 
resembles  the  displacement  of  hydrogen  by  zinc,  as  in  our  illustration. 
Most  changes  between  ionogens,  including  oxidations  and  double 
decompositions,  may  be  adapted  so  as  to  deliver  this  form  of  energy. 
It  need  hardly  be  added  that,  since  the  transformation  of  chemically 
equivalent  amounts  of  different  sets  of  substances  produces  very  dif- 
ferent quantities  of  heat,  so  it  produces  also  correspondingly  different 
amounts  of  electrical  energy.  Thus  the  original  free  chemical  energy 
may,  theoretically,  be  measured  by  either  method.  In  practice, 
however,  the  thermochemical  plan  fails  entirely  in  many  cases  (cf. 
pp.  27,  79),  and  the  electrical  is,  as  we  shall  see,  often  much  more 
instructive.  The  study  of  what,  to  parody  the  phraseology  of 
thermochemistry,  we  might  call  "  exoelectrical"  actions,  thus  resolves 
itself  into  constructing  experimental  battery-cells  involving  all  kinds 
of  chemical  changes,  and  studying  the  electric  currents  which  are  set 
in  motion  by  the  progress  of  the  changes.  We  have  therefore  named 
this  branch  of  the  science  electromotive  chemistry. 

In  addition  to  its  significance  theoretically,  electromotive  chem- 
istry has  recently  acquired  great  commercial  importance  because  of  the 
rapid  multiplication  of  electro-chemical  industries.  It  is  true  that 
the  majority  of  the  actions  used  in  these  industries  are  electrolytic 
(endoelectrical),  and  that  this  sort  of  change  is  the  precise  inverse  of 
the  other,  since  in  it  electricity  is  consumed  instead  of  set  in  motion, 

664 


ELECTROMOTIVE   CHEMISTRY  665 

but  it  is  also  true  that  neither  variety  can  be  understood  without  a 
study  of  both. 

Factors  and  Units  of  Electrical  Energy.  —  On  account  of  the 
close  relation  between  electromotive  chemistry  and  electrolysis,  parts 
of  the  former  subject  were  anticipated  when  the  latter  was  discussed 
(pp.  321-325).  These  pages  should  now  be  re-read  attentively.  In 
particular,  it  must  be  recalled  that  a  quantity  of  electrical  energy  is 
expressed  by  two  factors.  One  is  called  the  quantity  of  electricity,  and 
is  measured  in  coulombs.  The  other  is  called  the  electromotive  force 
in  the  case  of  a  current,  or,  when  a  current  is  not  flowing  or  is  not  be- 
ing considered,  the  difference  in  potential,  and  is  expressed  in  volts. 
Just  as  in  electrolysis  chemically  equivalent  quantities  of  elements  or 
ions,  in  being  liberated  from  solutions  of  different  substances,  use  up 
equal  quantities  of  electricity  (p.  317),  so  in  a  battery-cell  the  inter- 
action of  chemically  equivalent  amounts  of  different  sets  of  substances 
produces  equal  quantities  of  electricity  (p.  321).  Likewise,  just  as  in 
the  former  case  different  amounts  of  electrical  energy  (p.  323),  and 
therefore  different  electromotive  forces,  are  required  to  produce  in 
different  solutions  equivalent  amounts  of  chemical  change  (p.  324),  so 
in  the  latter  case  different  amounts  of  electrical  energy  are  generated 
by  the  complete  interaction  of  chemical  equivalents  of  different  sets  of 
substances,  and  therefore  diverse  differences  in  potential  are  created 
and  currents  of  different  electromotive  force  are  produced.  The  electri- 
cal energy  used  in  the  former  case  or  produced  in  the  latter  is  expressed 
by  the  product  of  the  factors  : 

No.  of  coulombs  x  No.  of  volts  =  Quant,  of  elect,  energy  (\njoules,  p.  25). 

If  we  consider  the  time  occupied  by  either  process,  and  wish  to  express 
the  rate  at  which  the  energy  is  consumed  or  produced,  we  regard  1 
coulomb  per  second  (1  ampere)  as  the  unit. 
Hence  : 

No.  of  amperes  x  No. )       (  Kate  of  production  or  consumption  of  elec- 
of  volts  )       (  trical  energy  (in  joules  per  sec.  =  watts). 

The  erg  (p.  25)  is  so  small  as  a  unit  of  energy  or  work,  that  the 
joule  (=  10,000,000  ergs)  and  the  kilojoule  (1000  joules)  are  more  often 
employed.  Similarly,  the  rate  at  which  the  energy  is  delivered  or  used 
(the  power)  is  expressed  by  the  watt  ( =  10,000,000  ergs  per  sec.)  or 
the  kilowatt  (1000  watts).  The  horse-power  is  736  watts.  Thus 


666  INORGANIC    CHEMISTRY 

the  electrolysis  of  one  formula-weight  of  hydrochloric  acid  (36.5  g.) 
requires  96,540  coulombs  (p.  323)  and  an  E.M.F.  of  1.41  volts  (p.  324). 
The  electrical  energy  needed  to  perform  this  work  is  therefore  96,540 
X  1.41  =  136, 121  joules. 

Displacement  Cells.  —  While  various  kinds  of  chemical  changes 
may  be  arranged  to  deliver  electricity,  we  shall  confine  our  attention 
almost  entirely  to  actions  in  which  one  free  substance  displaces 
another  from  combination.  Amongst  the  metals  the  order  of  displa- 
cing power  is  as  follows  (p.  362)  : 

Mg-Zn-Cd-Fe-Sn-Pb-H-Cu-Hg-Ag-Pt-Au. 

We  have  already  noted  many  of  the  single  facts  expressed  by  this  list. 
Thus,  metallic  copper  precipitates  metallic  mercury,  silver,  platinum, 
and  gold  from  solutions  which  contain  salts  of  these  elements,  but  it 
is  in  general  entirely  inactive  in  solutions  containing  salts  of  the 
elements  which  precede  it  in  the  list.  Conversely,  all  the  metals  pre- 
ceding copper  will  severally  precipitate  metallic  copper  from  a  solution 
of  a  cupric  salt,  but  the  metals  following  copper  do  not  affect  the 
solution  at  all.  Furthermore,  the  speed  with  which  the  precipitation 
is  effected  and  the  amount  of  heat  evolved  are  greater  when  copper 
precipitates. gold  than  when  it  precipitates  silver  :  they  are  also  greater 
when  magnesium  precipitates  copper  than  when  tin  does  so.  The 
same  statements  apply,  mutatis  mutandis,  to  each  of  the  other  ele- 
ments. This  is  the  behavior  when  the  experiment  is  made  in  the 
ordinary  way. 

One  preliminary  condition  is  indispensable  when  electricity  is  to 
be  obtained.  When  the  zinc  was  placed  in  sulphuric  acid  and  con- 
nected with  a  platinum  wire  (p.  20),  there  was  nothing  to  prevent  the 
metal  from  interacting  directly  with  the  acid,  and  generating  heat.  In 
fact,  it  did  so  act.  Hydrogen  appeared  on  the  zinc  itself,  as  well  as  on 
the  platinum,  and  both  the  heat  and  the  electricity  would  have  had  to 
be  counted  if  a  measurement  or  study  of  the  relation  between  the 
amount  of  zinc  and  acid  used  and  the  quantity  of  energy  produced  had 
been  in  question.  Since  this  is  precisely  the  present  problem,  all 
the  energy  must  now  be  secured  as  electricity,  and  the  only  way  to 
accomplish  this  is  to  prevent  the  mingling  of  the  interacting  materials. 
Paradoxical  as  it  may  seem,  it  is  easily  possible  to  get  the  electricity 
and  yet  fulfil  this  essential  condition.  The  plan  in  all  battery-cells  is 
to  place  the  one  substance  in  or  around  one  pole,  and  the  other  substance 


ELECTROMOTIVE   CHEMISTRY 


667 


around  or  in  the  other  pole,  and  to  separate  the  substances  by  a  porous 
wall  or  equivalent  arrangement.  The  figure  (Fig.  102)  shows  dia- 
grammatically  the  plan  that  will  serve  for  the  generation  of  electricity 
by  the  use  of  any  interaction. 

Suppose  the  action  is  a  displacement,  and  that,  for  example,  the 
metal  is  zinc  and  the  salt  cupric  sulphate  in  solution.  The  pole  on  the 
left  is  metallic  zinc ;  the  solution  to  the  right  of  the  porous  partition 
contains  the  cupric  sulphate.  To  complete  the  arrangement,  a  pole 
made  of  some  conductor  is  needed 
on  the  right,  and  a  conducting  solu- 
tion on  the  left.  For  these,  any 
substances  may  be  chosen,  provided 
only  that  they  do  not  interact  with 
the  adjacent  material.  If  they  do, 
part  of  the  two  main  substances 
will  be  used  up,  with  generation 
of  heat  instead  of  electricity.  For 
the  pole  we  may  take  copper  itself 
or  any  metal  following  it  in  the 
series  ;  and  copper  is  less  expensive 
than  any  of  these.  For  the  solution 
we  may  take  a  salt  of  zinc  or  of  any 

metal  preceding  zinc  in  the  series,  provided  that  the  salt  chosen  will 
not  interact  with  the  cupric  sulphate  which  it  meets  inside  the  porous 
wall.  Zinc  sulphate  or  sodium  chloride  will  serve  the  purpose.  The 
function  of  the  porous  partition  is  to  permit  the  migration  of  ions,  but 
prevent  a  general  mixing  of  the  materials.  It  is  evident  that,  when 
the  corresponding  conditions  are  observed,  any  other  pair  of  elements 
from  the  series  may  be  chosen  for  study.  In  practice,  it  is  common  to 
place  the  solution  of  the  substance  corresponding  to  the  cupric  sulphate 
along  with  its  electrode  in  an  outer  jar,  and  a  rod  of  the  more  active 
metal  along  with  its  conducting  fluid  in  an  inner  jar  made  of  porous 
earthenware.  The  following  are  the  phenomena  observed  : 

1.  The  pole  on  the  left  becomes  charged  negatively  on  account  of  the 
departure  of  positively  charged  ions  from  its  surface,  for  this  metal 
(zinc,  for  example)  goes  into  solution.  The  pole  on  the  right  becomes 
positively  charged  on  account  of  the  discharge  of  positive  ions,  and  de- 
position of  the  metal  (copper,  for  example)  on  its  surface.  A  current, 
therefore,  flows  from  +  to  —  through  the  wire,  and  passes  in  the  form 
of  migrating  ions  from  —  to  +  through  the  liquids.  The  negative 


FIG,  102. 


668  INORGANIC  CHEMISTRY 

ions  simultaneously  drift  away  from  the  pole  where  the  solution 
is  losing  metal,  towards  the  pole  where  it  is  gaining  metal,  and  so 
the  balance  of  the  two  sorts  of  ions  is  preserved  in  every  part  of  the 
liquids.  The  pole  on  the  left  is  the  anode,  that  on  the  right  the  cath- 
ode. This  whole  state  of  affairs  continues  until  either  all  the  copper 
is  precipitated,  or  all  the  zinc  is  consumed. 

2.  The  difference  in  potential  (E.M.F.)  is  not  affected  by  changes  in 
the  size  or  shape  of  the  poles  or  in  the  amounts  of  the  solutions,  provided 
the  materials  are  not  changed.     It  is  very  noticeably  affected,  however, 
by  alterations  in  the  concentrations  of  the  solutions.     In  particular, 
an  increase  in  the  concentration  of  the  ions  round  the  cathode  (the 
Cu"  ions,  for  example)  increases  the  difference  in  potential.     It  seems 
to  assist  the  action  by  helping  to  force  the  ions  out  of  the  solution. 

3.  In  general,  the  magnitude  of  the  electromotive  force  is  greater 
the  farther  the  chosen  metals  are  removed  from  one  another  in  the 
series.     Zinc  with  cupric  sulphate  gives  a  greater  difference  in  poten- 
tial than  cadmium  with  the  same  salt,  but  a  smaller  one  than  zinc  with 
silver  nitrate.     The  sum  of  the  differences  in  potential   of   the  pairs 
Zn-Cd",  Cd-Cu",  Cu-Ag*,  is  exactly  equal  to  the  difference  of  the  pair 
Zn-Ag',    provided   solutions   of  equivalent    ionic    concentration   are 
employed. 

4.  The  quantity  of  electricity  produced  depends  on  the  number  of 
equivalents  of  material  consumed.     The  strength  of  the  current  (the 
rate  at  which  the  electricity  is  developed)  depends  on  the  surface  of 
the  poles  and  other  circumstances  which  influence  the   rate  at  which 
the  materials  are  able  to  undergo  chemical  change.     One  gram-equiva- 
lent of  the  active  metal  (say,  32.7  g.  of  zinc)  in  becoming  ionized  will 
set  in  motion  96,540  coulombs,  and  deposit  one  gram-equivalent  of  the 
less  active  metal  (say,  31.8  g.  of  copper).     If  this  amount  were  to  be 
ionized  during  every  thirty  minutes  (  =  1800  seconds)  — and  plates  of 
great  surface  would  be  required  to  make  this  possible  —  the  current 
strength  would  be  96,540/1800  =  53.6  amperes. 

Potential  Differences  Produced  by  the  Metals  Singly.  — The 

facts  just  discussed  may  be  stated  in  various  ways.  For  example,  we 
may  say  that,  since  every  metal  (save  the  last  of  the  series)  displaces 
some  other  metal  or  metals,  each  metal  has  a  distinct  tendency  to 
become  ionic.  We  may  picture  this  as  a  pressure  or  tension  driving 
the  particles  into  solution,  and  actually  resulting  in  ionization  when 
the  necessary  electricity  is  available.  This  pressure  must  evidently 


ELECTROMOTIVE    CHEMISTRY  669 

be  greatest  with  magnesium,  and  least  with  gold.  On  the  other  hand, 
we  may  suppose  a  converse  tendency  of  the  ions  to  force  themselves 
out  of  solution  and  to  deposit  themselves  on  the  cathode,  this  tendency 
actually  becoming  continuously  operative  when  means  of  disposing  of 
the  electric  charges  is  provided.  This  tendency  is  evidently  greatest 
with  gold,  and  least  with  magnesium.  Now  the  use  of  cells  like  those 
described  above  does  not  enable  us  to  verify  these  hypotheses,  because 
in  each  cell  two  operations  are  going  on  simultaneously,  and  the  electri- 
cal effects  are  the  joint  result  of  both.  Fortunately  it  has  been  found 
possible  to  observe  and  measure  the  electrical  effects  at  each  end  of 
the  cell  separately.  It  is  found,  in  fact,  that  the  difference  in  poten- 
tial between  the  poles  is  for  the  most  part  made  up  of  the  differ- 
ences in  potential  between  each  pole  and  the  liquid  in  which  it  is 
immersed. 

After  the  cell  has  been  in  operation  for  some  time,  the  cathode 
system  consists  of  a  pole  covered  with  deposited  metal  (say  copper)  in 
contact  with  a  solution  containing  ions  of  this  same  element.  At  this 
end  the  ions  are  diminishing  in  number,  and  the  amount  of  free  metal 
is  increasing.  The  anode  system  likewise  consists  of  a  metal  (say 
zinc)  in  contact  with  a  solution  containing  ions  of  the  same  element. 
Here  the  number  of  ions  is  increasing,  and  the  metal  is  wearing  away. 
The  ions  of  the  foreign  salt,  if  such  a  salt  was  introduced  at  first,  need 
not  be  considered.  Thus,  for  the  purpose  of  making  a  strict  compari- 
son, we  may  immerse  each  metal  of  the  series  in  a  solution  of  one  of 
its  own  salts,  taking  the  latter  of  such  strength  that  there  is  normal 
concentration  of  the  metallion,  and  measure  the  difference  in  potential 
between  the  metal  and  the  solution. 

The  following  table  (p.  670)  contains  data  obtained  in  just  this  way. 
The  number  opposite  each  of  the  twenty-six  metals  (including  hydrogen) 
represents  the  potential  difference.  The  sign  preceding  the  number 
indicates  the  nature  of  the  electrical  charge  borne  by  the  solution. 

Two  or  three  examples  will  make  the  meaning  of  the  table  clearer. 
Opposite  Mg,  we  find  -f  1.21.  This  means  that  when  the  metal  mag- 
nesium is  immersed  in  a  solution  of  a  magnesium  salt  containing  a 
normal  concentration  of  Mg",  the  solution  becomes  positively  charged 
(the  metal,  negatively)  and  that  the  difference  in  potential  between 
metal  and  solution  is  1.21  volts.  Similarly  in  the  case  of  zinc,  the 
solution  is  positive,  but  the  potential  difference  is  smaller.  With 
copper  the  solution  is  negative  (and  the  metal  positive).  With  silver 
the  solution  is  negative,  and  more  strongly  so  than  with  copper.  Cobalt 


670 


INORGANIC   CHEMISTRY 


is  near  the  border  line,  the  normal  solution  of  cobaltion  being  very 
slightly  negative. 

POTENTIAL     DIFFERENCES     IN     VOLTS     FOR    NORMAL    SOLUTIONS 

OF   CATIONS. 


K 

(+2.9) 

Ni 

-0.04 

Na 

(  +  2.54) 

Sn  (Sn") 

-  0.08  ? 

Ba 

(  +  2.54) 

Pb 

-  0.13 

Sr 

(+2.49) 

H 

-0.28 

Ca 

(+  2.28) 

Cu  (Cu") 

-0.61 

Mg 

+  1.21 

As 

-  0.62  ? 

Al 

+  1.00 

Bi 

-0.67? 

Mn 

+  0.80 

Sb 

-0.74? 

Zn 

+  0.49 

Hg  (Hg") 

-1.03 

Cd 

+  0.14 

Ag 

-  1.05 

Fe  (Fe") 

+  0.06 

Pd 

-  1.07? 

Tl 

+  0.04 

Pt 

-1.14? 

Co 

-  0.04 

Au 

-  1.35? 

For  a  hydrogen  pole  a  piece  of  palladium  saturated  with  hydrogen 
gas  is  used.  The  values  for  the  metals  which  decompose  water  with 
ease  cannot  be  ascertained  by  direct  observation.  The  numbers  in 
parentheses  are  calculated  from  the  heats  of  ionization  and  serve 
simply  to  indicate  the  order  of  these  elements.  An  interrogation 
point  indicates  that  the  value  is  uncertain. 

These  facts  enable  us  to  state  in  more  definite  terms  the  formula- 
tive  hypothesis  foreshadowed  above.  It  was  first  put  forward  by 
Nernst. 

Every  metal  has  a  certain  solution  tension  or  pressure  tending  to 
drive  it  into  solution  (in  ionic  form,*  of  course,  since  it  is  not  soluble 
otherwise).  The  value  of  this  pressure  becomes  rapidly  less  as  we 
pass  through  the  series  from  magnesium  to  gold.  If  the  ions  of  the 
same  metal  are  already  present,  they  tend  to  give  up  their  electrical 
charges  and  deposit  themselves  upon  the  metal.  These  two 
tendencies  oppose  one  another  just  as  solution  pressure  and 
osmotic  pressure  oppose  one  another  in  the  ordinary  process 

*  The  idea  of  the  charge  of  electricity  is  apt  to  interfere  with  the  ready  accept- 
ance of  this  hypothesis.  If  it  is  remembered  that  the  ionic  form  of  an  element  is 
simply  an  allotropic  modification  (Ostwald)  with  a  different  amount  of  available 
energy,  the  difficulty  disappears.  In  the  ionic  allotrope  the  free  energy  is  some- 
times greater  (cobalt  to  gold)  and  sometimes  less  (potassium  to  thallium)  than  in 
the  free  element. 


ELECTROMOTIVE   CHEMISTRY  671 

of  dissolving  any  substance  (p.  152).  When  the  tendency  of 
the  ions  to  deposit  themselves  is  the  greater  of  the  two,  a  very  minute 
excess  of  deposition  over  solution  occurs,  and  thus  the  solution  has,  as 
a  whole,  a  negative  charge  (having  lost  some  positive  ions),  and  the 
metal  has  a  positive  charge  (having  acquired  it  from  the  deposit  of  a 
few  ions).  This  is  the  case  with  gold  and  the  metals  as  far  up  the 
list  as  cobalt.  When,  on  the  other  hand,  the  solution  pressure  of  the 
metal  is  the  greater  of  the  two,  the  solution  acquires  a  very  slight  excess 
of  positive  ions,  and  is,  therefore,  positively  charged  when  compared 
with  the  metal.  This  is  the  case  from  potassium  down  to  thallium. 

The  measure  of  the  "  tendency  of  the  ions  to  deposit  themselves  " 
is  simply  the  osmotic  pressure  of  the  metallions.  We  perceive  this  to 
be  the  case,  for,  when  we  take  a  stronger  solution  of  the  salt  and  there- 
fore an  increased  osmotic  pressure  of  the  ions,  an  instant  effect  is  pro- 
duced. The  solution  becomes  less  positive,  or  more  negative,  as  the 
case  may  be.  Evidently  the  solution  pressure  of  each  of  the  metals 
near  to  cobalt  is  almost  exactly  balanced  by  the  osmotic  pressure  of  a 
normal  solution  of  the  ions  composed  of  the  same  metal.  This  pressure, 
for  a  univalent  metal,  is  22.4  atmospheres  (p.  289),  and  for  a  bivalent 
metal  11.2  atmospheres.  The  metals  above  cobalt  have  solution  press- 
ures higher  and  higher  above  this  norm  ;  those  below  cobalt  have  solu- 
tion pressures  farther  and  farther  below  it.  The  effect  of  changing  the 
osmotic  pressure  is  independent  of  the  particular  substances  used,  and 
depends  only  on  the  valence.  When  the  concentration  of  the  metallion 
becomes  lO^V,  0.058  volts  must  be  subtracted  (algebraically)  from  the 
potential  (see  above  table)  of  the  liquid,  if  the  metallion  is  univalent. 
If  it  is  w-valeiit,  O.OoS/n  must  be  subtracted.  When  the  solution  is  0.1 
N,  0.058//1  volt  must  be  added  ;  when  it  is  0.01  N,  2  x  0.058/ra  must  be 
added,  and  so  forth.  Thus,  zinc  with  deci-normal  zincion  gives  -f-  0.49 
-f-  0.058/2  =  .52  volts,  approximately  ;  silver  with  centinormal  argent- 
ion  gives  —1.05  H-  (2  x  0.058)  =  —.93  volts,  approximately.  And,  in 
general,  if  c  be  the  equivalent  concentration  of  the  metallion  in  the 
liquid  under  consideration,  and  irc  the  electrical  potential  of  that  liquid, 
while  7nvr  is  the  potential  of  the  liquid  containing  N  metallion, 

0.058  .      1 

*C  =   TTN  +  -  -     lOg  -  . 

71  C 


Application  to  Cells.  —  When,  now,  a  cell  with  two  poles  and  two 
metallions  is  set  up,  we  can  tell  from  the  above  table  what  the  differ- 


672 


INORGANIC   CHEMISTRY 


ence  in  potential  between  the  two  poles  will  be.  We  may  regard  the 
two  systems — the  anodic  and  cathodic —  as  working  against  each  other. 
Each  metal  tends  to  project  its  ions  into  the  solution  and  to  generate 
a  positive  current  in  the  liquid  and  a  negative  one  in  the  wire.  If 
both  solutions  are  normal,  or,  in  general,  of  equal  equivalent  concen- 
tration, the  relative  solution  pressures  of  the  metals  decide  the  direc- 
tion of  the  resultant  current,  and  its  magnitude  will  be  the  difference 
of  the  two  effects.  Thus,  the  values  for  the  following  pairs  will  be  : 


Zn-Cd",  +0.49-  (  +  0.14)  =+0.35, 
Cd-Cu",  +  0.14-(-0.61)=+0.75, 
Zn-Cu", 


Zinc  the  negative  pole. 
Cadmium  the  negative  pole. 
Zinc  the  negative  pole. 

The  Daniell  or  gravity  cell  (Fig.  103)  represents  the  last  of  these 
three  combinations.  The  copper  pole  is  at  the  bottom,  and  the  zinc 

plate  is  suspended  above  it. 
The  cell  is  charged  with  a  dilute 
solution  of  sodium  chloride,  and 
blue  vitriol  crystals  are  thrown 
in  and  dissolve.  So  long  as 
the  contents  are  not  disturbed, 
the  solutions  require  no  porous 
septum  to  keep  them  apart.  It 
is  true  that,  when  the  current 
is  not  being  used,  and  the  cell 
is  not  working,  the  cupric  sul- 
phate diffuses  upwards.  During 
the  time  that  the  circuit  is 
closed,  however,  the  effects  of 

_--MM^<B__  diffusion   are  nullified  by   the 

migration  of  the  cuprion  away 

FlQ-  103'  from  the  zinc  and  towards  the 

positive  pole.  The  actual  elec- 
tromotive force  of  the  current  delivered  by  this  cell  is  a  little  over 
1  volt,  and  accords,  therefore,  with  the  value  calculated  from  the  poten- 
tial difference  observed  at  each  of  the  two  poles. 

The  cell  Zn-H'  (0.77  volts)  works  without  a  septum,  provided  the 
direct  action  of  the  zinc  on  the  acid  is  minimized  by  adequate  amalga- 
mation with  mercury.  It  gives  a  very  inconstant  electromotive  force, 
however,  because  the  platinum  plate  used  as  the  cathode  becomes  cov- 
ered with  bubbles  of  hydrogen,  and  so  the  internal  resistance  of  the  cell 


ELECTROMOTIVE   CHEMISTRY  673 

is  greatly  increased.  The  polarization  (p.  324)  also  diminishes  the 
electromotive  force.  These  difficulties  are  remedied,  and,  in  fact,  a 
great  increase  in  the  E.M.F.  of  the  cell  is  effected,  by  surrounding  the 
cathode  with  an  oxidizing  agent  which  shall  convert  the  hydrogen  into 
water.  The  energy  obtainable  is  thus  that  of  a  strong  oxidizing  agent 
on  zinc,  and  not  merely  that  of  an  acid.  In  the  Bunsen  cell  the  cath- 
ode is  a  carbon  block  surrounded  by  concentrated  nitric  acid.  In  the 
dichromate  battery  it  is  a  carbon  block  with  chromic  acid.  Each  of 
these  cells  gives  an  E.M.F.  of  1.9  volts.  In  the  Leclanche  cell  the  cath- 
ode is  a  mixture  of  carbon  and  manganese  dioxide,  and  the  fluid  is  a  so- 
lution of  ammonium  chloride  from  which  the  zinc  displaces  hydrogen. 
The  dioxide,  being  solid,  oxidizes  the  hydrogen  slowly,  and  the  cell  can 
be  used  for  only  a  few  minutes  at  a  time  without  becoming  polarized. 
The  E.M.F.  is  1.48  volts.  Dry  cells  are  of  the  same  nature,  but  con- 
tain a  porous  solid  which  holds  the  liquid  by  capillary  forces  (for 
Accumulators,  see  under  Lead). 

A  cell  is  thus  an  engine  for  the  direct  transformation  of  chemical 
into  electrical  energy,  just  as  a  steam-engine  transforms  chemical 
energy,  by  several  stages,  it  is  true,  into  mechanical  energy.  Accord- 
ing to  our  hypothesis  the  cell  is  driven  by  pressure-differences  in  the 
materials  in  and  around  the  two  poles. 

Other  Applications :  Couples:  Concentration  Cells.  —  We  are 

now  in  a  position  to  understand  the  effect  of  a  couple  (p.  96).  Zinc, 
in  contact  with  platinum  or  copper  and  immersed  in  acid,  is  practically 
a  short-circuited  cell,  and  it  is  found  that,  for  some  undetermined 
reason,  hydrogen  is  liberated  more  readily  from  the  surface  of  the 
platinum  or  copper  than  from  that  of  the  zinc.  Again,  galvanized  iron 
is  also  a  couple.  The  zinc  is  the  anode,  and,  when  dilute  carbonic  acid 
rests  on  a  damaged  part  of  the  surface  of  a  sheet  of  the  material,  and 
is  in  contact  with  both  metals,  the  zinc  is  ionized  and  passes  into 
combination  as  carbonate.  The  iron  is  the  cathode  and  is  not  affected. 
On  the  other  hand,  a  damaged  tin-plate  (tin  on  iron)  is  rapidly  rusted. 
Here  iron  is  the  anode,  and  goes  into  combination,  while  the  tin  is  the 
cathode.  The  ferrous  carbonate  (q.v.),  at  first  formed,  is  subsequently 
oxidized  and  gives  rust. 

When  two  strips  of  a  metal  are  immersed  together  in  a  solution  of  a 
salt  of  the  same  metallic  element,  —  for  example,  two  rods  of  tin  in  a 
normal  solution  of  stamious  chloride,  —  the  pieces  of  metal  show  no 
difference  in  potential  and  no  current  flows  when  they  are  connected 


674 


INORGANIC   CHEMISTRY 


FIG.  104. 


by  a  wire.     The  two  poles  and  their    solutions  are  here  alike  and 
—  .07  — (—.07)  =  0.*     But  if  the  solution  round  one   pole   is  diluted 

to  JV/10  concentration,  the  potential  at 
that  pole  becomes  at  once  —  .07  +  .058/2 
=  —  .04  volts,  approximately,  and  a  cur- 
rent is  set  up.  The  positive  current 
flows  from  the  pole  in  the  stronger  solu- 
tion through  the  wire  to  that  in  the 
dilute  solution,  and  thence  through  the 
liquid.  Thus,  tin  is  dissolved  from  the 
latter  pole  (the  anode),  and  the  concentra- 
tion of  the  solution  round  it  is  increased. 
On  the  other  hand,  tin  is  deposited  in 
long  needles  on  the  former  pole  (the  cath- 
ode), and  the  concentration  in  its  neigh- 
borhood is  correspondingly  decreased. 
The  figure  (Fig.  104)  shows  the  simplest 
arrangement,  where  the  more  concen- 
trated, denser  liquid  is  below,  and  one  rod 
of  tin,  passing  through  both  layers,  fur- 
nishes at  once  the  two  poles  and  the  con- 
nection. The  chloridion  migrates  through  the  solution,  in  a  direction 
opposite  to  that  taken  by  the  tin  ions,  and  thus  passes  upwards  into 
the  dilute  solution  to  balance  the  fresh  tin  ions  that  are  continuously 
formed.  All  change  ceases  when  the  concentrations  have  become 
equalized.  A  cell  of  this  kind  is  called  a  concentration  cell. 

The  concentration  cell  is  instructive  because  it  shows  that  the  order 
of  the  metals  in  the  electromotive  series  is  not  determined  by  the  metal 
alone,  but  also  by  the  concentration  of  the  solution.  The  order  of  the 
metals  in  the  electromotive  series  is  therefore  subject  to  variation.  An 
extreme  case  of  this  occurred  in  a  recent  chapter.  Zinc  displaces  cop- 
per from  a  solution  of  a  cupric  (or  cuprous)  salt,  and  any  but  a  prodi- 
giously dilute  solution  will  show  the  effect.  But  a  solution  containing 
cuprocyanion  Cu(CN)2'  has  precisely  this  very  minute  concentration  of 
copper  ions  which  will  turn  the  scale.  Hence,  zinc  will  not  displace 
copper  from  this  solution  (p.  624).  On  the  contrary,  copper  will  dis- 
place zinc  from  a  solution  of  a  salt  of  the  latter  containing  excess  of 
potassium  cyanide,  and  therefore  the  complex  salt  K.Zn(CN)s. 

*  In  copper  refining  (p.  617),  the  state  of  equilibrium  is  destroyed  and  the 
scale  turned  by  the  introduction  of  a  current  of  low  E.M.F.  from  a  dynamo. 


ELECTROMOTIVE   CHEMISTRY  675 

The  solubilities  of  insoluble  salts,  such  as  those  of  silver  (p.  630), 
have  also  been  measured  on  the  principle  of  the  above  tin  cell.  If  a 
saturated  solution  of  silver  chloride  is  placed  round  one  pole  of  silver, 
and  normal  argention  (say  from  silver  nitrate)  round  another,  a  con- 
centration cell  of  high  E.M.F.  is  formed.  The  magnitude  of  the 
difference  in  potential  enables  us  at  once  to  calculate  the  ratio  of  the 
concentrations  of  argention  round  each  pole^  and  therefore  to  get  at 
the  concentration  of  the  silver  chloride  solution. 

Electrolysis  :  Discharging  Potentials.  —  Another  kind  of  cell 
may  be  made  by  placing  a  rod  of  tin  (or  any  other  metal)  in  an  elec- 
trolyte on  one  side  of  a  porous  septum  (say  on  the  left  side  of  the  cell 
shown  in  Fig.  102),  and  chlorine  water  on  the  other.  An  indifferent 
ionogen  (say  KC1)  must  be  used  on  the  right  side  also,  because  chlorine 
water  is  not  a  good  conductor.  A  platinum  plate  is  likewise  required 
as  a  pole  at  this  end.  With  this  arrangement  the  tin  becomes  ionized 
and  renders  its  pole  negative.  Simultaneously  the  chlorine  in  contact 
with  the  platinum  becomes  ionic  and  leaves  the  other  pole  positive. 
Thus,  a  current  of  considerable  E.M.F.  flows  from  the  platinum  pole 
to  the  tin  pole  through  the  wire.  Tin  ions  and  chlorine  ions  are 
formed,  and  therefore,  potentially,  tin  chloride.  But  it  will  be  noted 
again  that  the  materials  have  to  be  kept  apart,  otherwise  direct  union 
and  mere  heat-production  will  result. 

Now  it  will  be  recalled  (p.  324)  that  when  a  salt  is  electrolyzed, 
the  materials  liberated  on  the  electrodes  generate  a  counter  current, 
called  the  polarization  current,  which  it  is  the  business  of  the  electro- 
lyzing  current  to  overcome.  It  will  now  be  clear  that  the  cell  just 
described  is  identical  with  the  arrangement  which  would  result  from 
electrolyzing  tin  chloride.  The  amount  by  which  the  E.M.F.  of  the 
electrolyzing  current  is  cut  down  (p.  324)  by  the  polarization  current 
is  exactly  equal  to  the  potential  of  a  cell  like  the  above  containing  the 
same  elements.  The  minimum  difference  in  potential  which  will  de- 
compose an  electrolyte  is  called  the  discharging  potential  and  is  nu- 
merically equal  to  the  E.M.F.  of  the  corresponding  battery-cell,  but  of 
opposite  sign.  Furthermore,  since  the  potential  of  a  cell  is  made  up 
of  two  simple  potentials,  the  discharging  potential  must  be  made  up  of 
the  same  two  (discharging)  potentials.  Thus,  with  the  help  of  a  list 
of  anions  and  their  single  potentials,  we  can  calculate  at  once  the  total 
E.M.F.  required  to  electrolyze  any  salt.  A  few  of  the  values  are  as 
follows : 


676  INORGANIC   CHEMISTRY 

POTENTIAL   DIFFERENCES   FOR   ANIONS   (IN   VOLTS). 


I 

-  0.80 

OH 

-  1.96 

Br 

-  1.27 

S04 

-2.2 

O 

-  1.36 

HS04 

-  2.9 

Cl 

-1.69 

As  before,  the  anode  potential  is  supposed  to  work  against  the  cathode 
potential  and  is  subtracted  from  it.  Thus,  the  tin-chlorine  cell  pro- 
duces —  0.08  — (—1.69)  =  1.61  volts,  and  this  E.M.E.  will  just  suffice 
to  electrolyze  tin  chloride.  Similarly,  hydrochloric  acid  will  require 
at  least  -  0.28-(-  1.69)  ==  1.41  volts,  zinc  sulphate  0.49  -(-  2.2) 
=  2.69  volts. 

Oxygen  acids  like  sulphuric  acid  show  a  trace  of  decomposition  at 
1.08  volts  ( =  —  0.28 —  ( — 1.36)),  and  a  noticeable  but  still  small  decom- 
position at  1.68  volts  (=  -0.28  — (  —  1.96)),  due  to  the  H*  and  0"  and 
the  H*  and  OH'  respectively.  But  it  is  only  when  the  E.M.F.  reaches 
the  values  for  H*  and  SO/',  and  H'  and  HSO/,  namely,  1.92  and 
2.62  volts,  that  rapid  electrolysis  begins.  This  observation  answers, 
incidentally,  the  question  whether  in  the  so-called  "  electrolysis  of 
water,"  when  dilute  sulphuric  acid  is  used,  it  is  the  water  or  the  acid 
that  is  decomposed.  The  H"  and  OH'  decomposition  at  1.68  volts  is 
very  slight,  because  of  the  small  concentration  of  the  OH',  and  a  lens 
is  required  for  its  recognition.  The  more  vigorous  action  resulting 
from  the  discharge  of  S04"  and  HSO/  by  the  use  of  2-3  volts  is 
therefore  the  one  invariably  used. 

In  view  of  the  foregoing  facts,  it  is  probably  most  correct  to  say 
that  when  dilute  sulphuric  acid  is  electrolyzed,  e.g.  as  a  lecture 
experiment,  the  oxygen  liberated  at  the  anode  comes  mainly  from  a 
secondary  interaction  of  the  discharged  material  of  the  anions  with 
the  water  (p.  95).  A  minute  proportion  of  the  oxygen  in  such  an 
experiment  does  arise  from  primary  electrolysis  of  the  water,  but  this 
effect  of  the  current  is  in  itself  too  slight  to  be  visible  at  a  distance. 
When,  on  the  other  hand,  the  solution  electrolyzed  contains  a  salt  of 
sodium,  and  hydrogen  is  liberated  at  the  cathode,  this  gas  must  prob- 
ably be  regarded  as  coming  chiefly  from  primary  electrolysis  of  the 
water.  The  discharging  potential  for  sodium  chloride  should  be  -j-  2.54 
—  (—1.69)  =  4.23  volts,  and  with  the  help  of  a  mercury  cathode  a 
sodium  amalgam  is  easily  obtained  (p.  554).  But  it  will  be  found 
that,  with  platinum  electrodes,  hydrogen  and  chlorine  are  liberated 
freely  from  a  solution  of  salt  by  a  current  of  little  more  than  half  the 
above  mentioned  E.M.F.  The  positive  electricity  is  carried  in  the 


ELECTROMOTIVE   CHEMISTRY  677 

liquid  mainly  by  the  very  numerous  sodium  ions.  But,  apparently, 
when  these  ions  reach  the  cathode,  the  potential  difference,  being 
insufficient  to  discharge  the  natrion,  liberates  the  hydrion  of  the  water 
instead.  Thus  the  accumulating  hydroxidion  of  the  water,  and  the 
natrion  arriving  by  migration,  together  constitute  the  sodium  hydroxide 
which  is  another  product  of  this  electrolysis.  With  higher  E.M.F. 
and  sufficient  current  density,  natrion  is  doubtless  actually  discharged, 
and  in  that  case  a  part  of  the  hydrogen  liberated  is  furnished  by  the 
interaction  of  the  metal  with  the  water. 

The  ordinary  chemical  behavior  of  the  halogens  accords  with  the 
order  of  their  potential  differences.  Bromine  displaces  iodine,  and 
chlorine  displaces  both  (p.  361).  Chlorine,  however,  does  not  displace 
easily  perceptible  amounts  of  oxygen  from  water,  because  of  the  small 
concentration  of  the  0"  (obtained  by  secondary  ionization  of  the 
OH').  The  oxygen  freely  liberated  in  sunlight  comes  from  the  de- 
composition of  the  HC10  (p.  269).  Fluorine,  however,  which  would 
probably  show  a  potential  difference  below  that  of  the  much  more 
plentiful  OH',  displaces  oxygen  vigorously  by  discharging  this  ion. 

In  electrolyzing  a  mixed  solution  with  a  moderate  current,  the 
cations  and  anions  with  the  lowest  discharging  potentials  are  first 
liberated.  Thus,  silver  (—  1.05)  appears  before  copper  (—  0.6). 
Hence,  in  copper  refining  (p.  617),  the  copper,  of  which  there  is  a 
continuous  supply,  is  deposited,  and  the  more  active  metals  remain 
combined.  Indeed,  the  E.M.F.  used  is  not  sufficient  in  any  case  to 
discharge  them. 

The  Factors  of  Energy.  —  We  have  seen  that  the  amount  of  a 
given  supply  of  electrical  energy  is  described  by  two  factors,  the 
E.M.F.  and  the  quantity  of  electricity,  and  that  the  weight  of  material, 
which,  by  its  influence,  undergoes  a  given  chemical  change,  is  propor- 
tional solely  to  this  second  factor.  On  the  other  hand,  the  question 
whether  the  supply  of  energy  can  initiate  the  change  at  all  depends  on 
the  magnitude  of  the  first  factor  alone  (p.  324).  The  total  amount  of 
available  energy  does  not  influence  the  result  if  the  E.M.F.  is  not  above 
a  certain  minimum,  which  differs  from  case  to  case.  Now  the  same  is 
true  of  other  kinds  of  energy.  The  quantity  of  each  may  be  expressed 
as  the  product  of  an  intensity  factor  and  a  capacity  factor.  The  mag- 
nitude of  the  former  determines  whether  the  energy  can  be  trans- 
ferred or  transformed  or  not.  Heat  energy,  no  matter  how  much  of  it 
is  at  hand,  can  neither  flow  nor  be  transi'ormed  into  work  unless  the 


678  INORGANIC   CHEMISTRY 

source  is  at  a  higher  temperature  than  the  surroundings.  A  head  of 
water  will  do  work  only  when  it  is  connected  with  a  receptacle  at  a 
lower  level.  It  is  the  pressure  of  the  water  that  determines  its  avail- 
ability. The  E.M.F.  is  the  corresponding  factor  of  electrical  energy. 

Now  we  may  presume  that  chemical  energy  can  be  expressed  by 
two  factors.  One  of  these,  the  capacity  factor,  must  be  proportional 
to  the  quantity  of  material,  in  other  words,  to  the  number  of  chemical 
equivalents.  The  other  is  the  chemical  potential.  A  chemical  change 
which  does  not  take  place  on  a  small  scale  will  not  take  place  when 
more  material  is  used,  provided  the  relative  amounts  of  the  interact- 
ing substances  and  the  conditions  remain  unchanged.*  We  have,  in 
fact,  been  assuming  all  along  that  this,  the  capacity  factor,  is  not 
the  most  significant  one.  But  we  have  devoted  ourselves  to  noting 
such  things  as  these  :  that  chlorine  will  displace  bromine,  and  there- 
fore has  the  higher  potential  of  chemical  energy;  that  magnesium 
reduces  sand,  while  hydrogen  does  not,  and  that  magnesium  is  therefore 
a  more  active  reducing  agent ;  and  that  hypochlorous  acid  will  oxidize 
indigo,  while  free  oxygen  will  not,  and  is  therefore  a  more  powerful  oxi- 
dizing agent.  When  we  were  comparing  degrees  of  activity,  therefore, 
we  were  really  trying  to  describe  the  relative  potential  of  the  chemical 
energy  in  all  sorts  of  substances.  At  present  the  state  of  the  science 
permits  this  to  be  done  in  most  cases  in  a  rough  fashion  only. 

Since  the  capacity  factor  of  chemical  energy  is  proportional  to  the 
number  of  equivalent  weights  transformed,  and  the  capacity  factor  of 
electrical  energy  is  proportional  to  the  same  thing  (Faraday's  law),  it 
follows  that  the  intensity  factor  of  the  chemical  energy  (the  chemical 
potential)  in  a  given  substance  undergoing  a  given  change,  must  be 
proportional  to  the  corresponding  factor  (the  E.M.F.)  of  the  electrical 
energy  produced  when  the  same  change  takes  place  in  a  suitable  cell. 
The  potential  differences  described  above  are  therefore  often  much  more 
significant  than  are  the  results  of  thermochemical  measurements,  for 
the  latter  attempt  to  give  only  the  gross  quantity  of  chemical  energy 
(in  terms  of  the  equivalent  amount  of  heat  energy),  and  not  the  values 
of  the  factors.  The  potential  differences  come  nearer,  therefore,  to 
giving  us  absolute  values  for  chemical  activity  than  do  any  other  data 
we  possess. 

As  we  have  noted  before  (p.  78),  in  spite  of  the  enormous  range 
of  temperature  at  our  disposal,  extending  to  a  point  far  above  2500°  in 

*  Change  in  concentration,  however,  does  affect  activity,  and  therefore  modi- 
fies the  chemical  potential. 


ELECTROMOTIVE   CHEMISTRY  679 

the  electric  furnace,  there  are  many  substances  for  whose  decomposition 
a  sufficient  potential  of  heat  energy  is  not  available.  On  the  other 
hand,  amongst  substances  that  are  capable  of  furnishing  an  electro- 
lyte, when  dissolved  in  a  suitable  solvent  or  when  fused,  there  are  few 
that  are  not  decomposable  by  a  current  with  an  E.M.F.  of  less  than  10 
volts.  Hence  even  the  elements  which  give  the  most  stable  compounds 
and  are  the  most  difficult  to  isolate,  such  as  calcium  and  aluminium, 
are  liberated  by  electrical  methods  with  extreme  ease. 

Methods  of  Measuring  Chemical  Activity.  —  The  following  is 
a  summary  of  the  methods  of  measuring  chemical  activity. 

The  thermochemical  method  (p.  79)  can  be  used  in  every  chemi- 
cal change.  But  the  heats  of  reaction  represent  the  free  energy,  and, 
therefore  the  affinity,  only  when  the  heat  capacity  of  the  products  is 
equal  to  that  of  the  factors  and  no  changes  in  concentration  arise. 

For  measuring  the  activity  of  acids  in  dilute  solution,  several 
methods  have  been  mentioned  :  The  speed  of  interaction  of  different 
acids  with  the  same  metal  (p.  347)  ;  the  acceleration  of  the  speed  of 
hydrolysis  of  ethyl  acetate  (p.  504)  and  of  cane-sugar  (p.  500)  by 
different  acids  ;  the  amounts  of  insoluble  salts,  such  as  calcium  oxa- 
late  (p.  598),  or  zinc  sulphide,  which,  when  the  system  has  reached 
equilibrium,  are  found  to  have  been  decomposed  by  different  acids  under 
like  conditions ;  the  relative  extents  of  the  hydrolysis  of  salts  of  differ- 
ent weak  acids  (p.  344) ;  the  electrical  conductivity  (p.  325)  and  the 
freezing-  and  boiling-points  of  solutions  of  acids  (pp.  292,  293).  These 
last  measure  by  physical  methods  the  same  thing  that  the  others  de- 
termine by  chemical  means,  namely,  the  tendency  to  ionization  on 
which  the  activity  of  acids  depends  (p.  347.  See  also  p.  356). 

For  measuring  the  activity  of  bases,  we  have  :  The  relative  speeds 
of  saponification  of  esters  by  different  bases  (p.  505) ;  the  relative  ex- 
tents of  the  hydrolysis  of  salts  of  different  weak  bases  (p.  344)  ;  the 
conductivity  and  the  freezing-  and  boiling-point  methods,  which  meas- 
ure by  physical  means  the  tendency  to  ionization. 

For  measuring  the  relative  activities  of  metals  and  non-metals, 
we  have  :  The  single  potential  differences  (pp.  670,  676)  ;  and,  for  the 
former,  the  speed  of  interaction  of  different  metals  with  the  same  acid 
(p.  111). 

For  measuring  the  relative  activity  in  non-reversible  actions,  we 
have :  The  speed  with  which  the  actions  take  place  under  like  con- 
ditions (p.  250). 


680  INORGANIC   CHEMISTRY 

For  measuring  the  relative  activities  of  the  opposed  actions  in 
reversible  changes,  we  have:  The  concentrations  of  the  materials  re- 
maining when  equilibrium  has  been  reached  (p.  254).  The  relative 
activities  in  different  reversible  changes  may  also  be  ascertained  by 
comparing  the  concentrations  in  one,  at  equilibrium,  with  those  in 
another  (cf.  p.  257). 

For  measuring  the  relative  activities  of  oxidizing  and  reducing 
agents,  we  have  :  The  potential  differences  in  cells  arranged  after  the 
manner  of  the  Bunsen  and  Leclanche  cells  (p.  673). 

If  we  consider  the  whole  mass  of  phenomena,  it  must  be  admitted 
that  the  scientific  study  of  the  quantities  of  material  has  reached  a  far 
higher  level  of  exactness,  and  has  very  much  more  nearly  enveloped 
the  whole  field  covered  by  the  science,  than  has  the  study  of  relative 
activity.  Yet  it  is  evident  that  within  the  past  few  years  substantial 
advances  have  been  made  in  this  direction  also. 

Exercises.  —  1.  What  will  be  the  E.M.F.  of  each  of  the  following 
cells  when  each  of  the  metallions  is  present  in  normal  concentration : 
Mn  — Cu",  Cd  —  Pb"? 

2.  What  will  be  the  E.M.F.  of  a  concentration  cell  in  which  the 
poles  are  of  lead  and  the  plumbion  is  one  hundred  times  more  concen- 
trated round  one  pole  than  round  the  other  ? 

3.  What  will  be  the  discharging  potential   for  solutions  of  the 
following  substances,  if  we  assume  that  the  concentration  of  the  ions  is 
normal :  manganous  chloride,  hydrogen  bromide  ? 

4.  What  weight  of  aluminium  must  become  ionized  every  hour  in 
a  cell  in  order  that  a  current  of  five  amperes  strength  may  be  produced  ? 
What  would  be  the  E.M.F.  of  the  current  if  an  acid  with  normal  con- 
centration of  hydrion  surrounded  the  cathode  and  a  solution  of  normal 
aluminion   the   anode  ?     How  would   this  E.M.F.  be  affected  if  the 
aluminion  was  only  one-hundredth  normal  ? 


CHAPTER   XXXIX 
ALUMINIUM   AND    THE   METALS    OF   THE   EARTHS 

THE  fourth  column  of  the  periodic  table  (p.  411)  contains  boron  and 
aluminium  along  with  a  number  of  rare  elements.  The  chief  members 
of  the  family  are :  boron  (B,  at.  wt.  11),  aluminium  (Al,  at.  wt.  27. 1), 
gallium  (Ga,  at.  wt.  70),  indium  (In,  at.  wt.  115),  thallium  (Tl,  at.  wt. 
204.1),  all  on  the  right  side  of  the  column,  and  scandium  (Sc,  at.  wt. 
44.1),  yttrium  (Yt,  at.  wt.  89),  lanthanum  (La,  at.  wt.  138.9),  samarium- 
(Sa,  at.  wt.  150.3),  and  ytterbium  (Yb,  at.  wt.  173)  on  the  left  side. 
These  elements  are  all  trivalent. 

The  Hare  Elements  of  this  Family.  —  The  oxide  and  hydroxide 
of  boron  are  acidic  (p.  527).  Those  of  aluminium  (A1(OH)3),  gallium 
(Ga(OH)3),  indium  (In(OH)3),  and  thallium  (T10.0H)  are  basic,  but 
behave  also  as  acids  towards  strong  bases. 

Gallium  and  indium  occur  occasionally  in  zinc-blende,  and  were 
discovered  by  the  use  of  the  spectroscope.  The  former  takes  its 
name  from  the  country  (France)  in  which  the  discovery  was  made,  and 
the  latter  from  two  blue  lines  shown  by  its  spectrum.  Indium  gives 
a  complete  series  of  compounds  in  which  it  is  trivalent,  and  the 
chlorides  InCl  and  InClg  are  also  known. 

Thallium  is  found  in  some  specimens  of  pyrite  and  blende.  It  was 
discovered  by  Crookes  by  means  of  the  spectroscope  in  the  seleniferous 
deposit  from  the  flues  of  a  sulphuric  acid  factory.  It  received  its 
name  from  the  prominent  green  line  in  its  spectrum  (Gk.  0aAAos,  a 
green  twig).  It  gives  two  complete  series  of  compounds.  In  those  in 
which  it  is  trivalent  (thallic  salts),  it  resembles  aluminium  (q.v.).  Thus, 
the  salts  of  this  series  are  more  or  less  hydrolyzed  by  water.  Univalent 
thallium  recalls  both  sodium  and  silver.  Thallous  hydroxide  (T10H) 
is  soluble,  and  gives  a  strongly  alkaline  solution.  The  chloride  is 
insoluble  in  cold  water.  The  solutions  of  the  thallous  salts  are  neutral. 
The  metal  is  displaced  from  its  salts  by  zinc. 

Of  the  elements  on  the  left  side  of  the  column,  scandium,  whose 
existence  and  properties  were  predicted  by  Mendelejeff  (p.  412),  is  the 

681 


682  INORGANIC  CHEMISTRY 

best  known.  The  nietals  of  the  rare  earths,  of  which  it  is  one,  are 
found  in  rare  minerals  such  as  euxenite,  gadolinite,  orthite,  and  mona- 
zite,  which  occur  in  Sweden,  Greenland,  and  the  United  States. 
Cerium  (Ce,  at.  wt.  140.25),  neodymium  (Nd,  at.  wt.  143.6),  and  pras- 
eodymium (Pr,  at.  wt.  140.5),  occur  along  with  lanthanum  in  cerite,  a 
silicate  of  these  four  elements.  These  four  are  included  amongst  the 
metals  of  the  rare  earths.  The  compounds  of  many  of  these  rare  ele- 
ments behave  so  much  alike  that  separation  is  difficult.  It  is  certain, 
however,  that  there  are  several  with  atomic  weights  near  to  that  of 
lanthanum  for  which  accommodation  cannot  easily  be  found  in  the 
periodic  table,  and  some  of  these  are  probably  mixtures  of  still  more 
closely  related  elements.  Ostwald  has  compared  them  to  a  group  of 
minor  planets  such  as  in  the  solar  system  takes  the  place  of  one  large 
planet. 

ALUMINIUM. 

The  Chemical  Relations  of  the  Element.  —  Aluminium  is  triva- 
lent  exclusively.  Its  hydroxide,  like  that  of  zinc  (p.  648),  is  feebly 
acidic  as  well  as  basic,  and  hence  the  metal  forms  two  sets  of  compounds 
of  the  types  Nag.AlOg  and  A12.(S04)8.  The  salts  of  both  series  are  more 
or  less  hydrolyzed  by  water,  the  former  very  conspicuously  so.  It  is 
worth  noting  that  the  hydroxides  of  the  trivalent  metals,  or  metals  in 
the  trivalent  condition,  such  as  A1(OH)8,  Cr(OH)3,  Fe(OH)3,  are  all 
distinctly  less  basic  than  are  those  of  the  bivalent  metals  such  as 
Zn(OH)2,  Cd(OH)2,  Fe(OH)2,  Mn(OH)2.  This  fact  is  used  in  analysis 
(cf.  also  p.  661)  in  separating  the  two  sets.  When  a  solution  of  the 
chlorides  is  shaken  with  precipitated  barium  carbonate,  the  free  acid 
from  the  more  highly  hydrolyzed  salts  of  Al"",  Cr*"  and  Fe""  inter- 
acts with  this  substance,  the  hydrolysis  is  promoted : 

A1C13  +  3H20  <=>  A1(OH)3  +  3HC1, 

and  eventually  the  hydroxides  A1(OH)8,  Cr(OH)3,  and  Fe(OH)3  are 
completely  precipitated.  The  chlorides  of  the  bivalent  metals  remain 
in  the  solution.  Aluminium  does  not  enter  into  complex  anions  or 
cations,  and  is  too  feebly  base-forming  to  give  salts  like  the  carbonate 
or  sulphite. 

Occurrence.  —  Aluminium  is  found  very  plentifully  in  combina- 
ion,  coming  next  to  oxygen  and  silicon  in  this  respect.  The  feldspars 
^such  as  KAlSi308),  the  micas  (such  as  KAlSi04),  and  kaolin  (clay 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS     683 

H2Al2(Si04)2,H20),  are  the  commonest  minerals  containing  it.  Gar- 
nets, which  are  found  in  metamorphic  rocks,  are  mainly  an  orthosili- 
cate  of  calcium  and  aluminium  Ca3Al2(Si04)3.  Turquoise  is  a  hydrated 
phosphate  A12(OH)3P04,H20,  and  cryolite  a  double  fluoride  3NaF,AlF8. 
Various  forms  of  the  oxide  and  hydroxide  (see  below)  are  also  not 
uncommon  minerals. 


Preparation  and  Physical  Properties.  —  The  metal  is  now 
made  on  a  large  scale  by  electrolysis  of  the  oxide  (A1203)  dissolved  in 
a  bath  of  molten  cryolite.  The  operation  is  conducted  in  cells,  the 
carbon  linings  of  which  form  the  cathodes.  The  anodes  are  rods  of 
carbon  which  combine  with  the  oxygen  as  it  is  liberated.  The  metal 
sinks  to  the  bottom  of  the  cell  and  is  drawn  off  periodically,  while  fresh 
portions  of  the  oxide  are  added  from  time  to  time.  A  current  density 
of  5  amperes  per  sq.  cm.  of  cathode  area  and  an  E.M.F.  of  5-6  volts 
maintain  the  temperature  of  the  molten  materials,  and  cause  the 
decomposition. 

The  metal  melts  at  600-700°,  but  is  not  mobile  enough  to  make 
castings.  It  is  exceedingly  light  (sp.  gr.  2.6),  and  in  hardness  and 
tensile  strength  is  the  equal  of  any  of  the  other  metals,  with  the 
exception  of  steel.  It  has  a  silvery  luster,  and  scarcely  tarnishes,  the 
firmly  adhering  film  of  oxide  first  formed  protecting  its  surface. 
Although,  comparing  cross-sections,  it  is  not  so  good  a  conductor  of 
electricity  as  is  copper,  yet  weight  for  weight  it  conducts  better.  It  is 
difficult  to  work  on  the  lathe  or  to  polish,  because  it  sticks  to  the  tools, 
but  the  alloy  with  magnesium  (6-30  per  cent)  called  magnalium  has 
admirable  qualities  in  these  respects.  Aluminium  bronze  (5-12  per 
cent  aluminium)  is  easily  fusible,  has  a  magnificent  golden  luster,  and 
possesses  mechanical  and  chemical  resistance  exceeding  that  of  any 
other  bronze. 

The  metal  and  its  alloys  are  used  for  making  cameras,  opera- 
glasses,  cooking  utensils,  and  other  articles  requiring  lightness  and 
strength.  The  powdered  metal,  mixed  with  oil,  is  used  in  making  a 
silvery  paint. 


Chemical  Properties.  —  The  metal  displaces  hydrogen  from 
hydrochloric  acid  very  easily.  In  sulphuric  and  nitric  acid,  however, 
it  receives  a  coating  of  the  hydroxide,  formed  by  hydrolysis  of  the 
salt,  and  the  action  is  slow  in  the  former  case,  and  almost  nil  in 


684  INORGANIC  CHEMISTRY 

the  latter.  It  displaces  hydrogen  also  from  boiling  solutions  of 
the  alkalies,  forming  aluminates  : 

2A1  +  6NaOH  -»  2Na3A103  +  3H2. 

• 

In  consequence  of  its  very  great  affinity  for  oxygen,  aluminium  dis- 
places all  the  metals,  save  magnesium,  from  their  oxides.  Thus,  when 
a  mixture  of  aluminium  powder  and  ferric  oxide  is  placed  in  a  crucible 
and  ignited  by  means  of  a  piece  of  burning  magnesium  ribbon, 
aluminium  oxide  and  iron  are  formed  : 


Fe203  -f  2A1  -»  ALjOg  +  2Fe. 

The  very  high  temperature  (about  3000°)  produced  by  the  action  is 
sufficient  to  melt  both  the  iron  (m.-p.  1700°)  and  the  oxide  of  alumin- 
ium. The  products,  not  being  miscible,  separate  into  two  layers.  This 
very  simple  method  of  making  pure  specimens  of  metals  like  chromium, 
uranium,  and  manganese,  whose  oxides  are  otherwise  hard  to  reduce,  is 
called  by  Groldschmidt,  the  inventor,  "  aluminothermy."  The  sul- 
phides, such  as  pyrite,  are  reduced  with  equal  vigor  by  aluminium. 

Aluminium  Chloride.  —  If  the  metal  or  the  hydroxide  is  treated 
with  hydrochloric  acid,  and  the  solution  is  allowed  to  evaporate,  crys- 
tals of  A1C13,6H2O  are  formed.  When  heated,  this  hydrate  is  com- 
pletely hydrolyzed,  hydrochloric  acid  is  given  off,  and  only  the  oxide 
remains.  The  anhydrous  chloride  is  much  used  as  a  catalytic  agent 
for  causing  combination  in  organic  chemistry.  It  is  made  by  passing 
dry  chlorine  over  aluminium,  or  by  heating  the  oxide  with  carbon  in  a 
stream  of  chlorine.  Just  as  in  the  case  of  silicon  dioxide  (p.  520), 
neither  carbon  nor  chlorine  alone  will  act  upon  the  oxide. 

Aluminium  chloride  gives  a  vapor  pressure  of  760  mm.  at  183°,  and 
sublimes,  as  a  white  crystalline  solid,  without  melting.  Under  pres- 
sure, it  melts  at  193°.  In  the  mode  of  preparation  described  above,  it 
is,  therefore,  vaporized,  and  condenses  in  a  cool  part  of  the  tube.  It 
fumes  when  exposed  to  moist  air  on  account  of  the  hydrogen  chloride 
produced  by  hydrolysis,  and  only  with  excess  of  hydrochloric  acid  does 
it  give  a  clear  solution  free  from  basic  salts. 

Aluminium  Hydroxide  and  the  Aluminates.  —  When  an  alkali 
is  added  to  a  solution  of  a  salt  of  aluminium,  the  hydroxide  A1(OH)3  is 
precipitated  in  gelatinous  form.  It  is  a  white  hydrogele  (p.  523),  and 
loses  water  gradually  when  dried,  without  forming  any  intermediate 
hydroxides,  until  A1208  alone  remains.  It  interacts  both  with  acids 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS     685 

and  with  bases,  and  is,  therefore,  like  zinc  hydroxide  (p.  648),  ionized 
both  as  a  base  and  as  an  acid.  It  interacts  only  slightly  with  ammonium 
hydroxide,  because  this  substance  is  too  feebly  basic,  but,  from  the  solu- 
tion in  the  active  alkalies,  the  aluminates  Na8.A103,  Na.  A102,  and*  K.A102, 
can  be  obtained  in  solid  form .  Natural  forms  of  this  substance  are  hy  drar- 
gyllite  A1(OH)8  (=  A1203,3H20),  bauxite  A120(OH)4  (=  A^08,2H20), 
which  always  contains  ferric  oxide,  and  diaspore  A10.0H  (=  A1208, 
K,0). 

Commercially,  the  hydroxide  is  made  by  heating  bauxite  with 
sodium  carbonate.  The  ferric  oxide,  having  no  tendency  to  form  a 
carbonate  or  to  interact  with  a  base,  remains  unchanged.  The  sodium 
aluminate  which  is  formed  can  be  extracted  with  water : 

A120(OH)4  4  Na2C03  ->  2NaA102  4  C02  4-  2H20. 

The  hydroxide  is  then  precipitated  by  passing  carbon  dioxide  through 
the  solution : 

2NaA102  4-  C02  4-  3H20  ->  Na2CO8  +  2A1(OH)8. 
The  aluminates  are  largely  hydrolyzed  by  water  : 
NaAlO2  4  2H2O  +±  NaOH  4  A1(OH)8. 

Hence  an  excess  of  sodium  hydroxide  is  required  for  the  complete 
solution  of  aluminium  hydroxide  by  the  reversal  of  this  action.  Sodium 
aluminate  is  used  as  a  mordant  in  dyeing  (see  below),  on  account  of  the 
ease  with  which  the  solution  gives  up  aluminium  hydroxide  when  any 
material  is  present  which  can  combine  with  the  free  portion  of  the 
hydroxide  and  so  cause  forward  displacement  of  the  above  equilibrium. 
When  calcium  chloride  is  added  to  a  solution  of  sodium  aluminate, 
the  insoluble  calcium  metaluminate  is  deposited : 

2NaA102  4  CaCl2  -» Ca(A102)2  4-  2NaCl. 

The  relations  of  these  various  substances  are  shown  by  the  following 
formulae  : 


/0-H        /Q-Na  000 

Al-O-H     Al-O-Na    AT?  Al^ 

^0-H         ^0-Na         X°-H          N°-*a     A1/0 


686  INORGANIC   CHEMISTRY 

A  number  of  insoluble  metalumiDates  are  found  in  nature.  They 
crystallize  in  the  regular  system,  and  are  known  as  spinelles.  They 
contain  bivalent  metals  in  place  of  the  calcium  in  the  last-named 
compound.  Thus  we  have  spinelle  proper  Mg(A102)2,  and  gahnite 
Zn(A102)2.  Corresponding  and  isomorphous  derivatives  of  chromic 
and  ferric  hydroxides  are  chromite  Fe(Cr02)2  and  magnetite  Fe(Fe02)2. 

Aluminium  Oxide.  —  The  oxide  (alumina)  is  found  in  nature  in 
pure  form  as  corundum.  This  mineral  is  only  one  degree  less  hard 
than  the  diamond.  Emery  is  a  common  variety,  contaminated  with 
ferric  oxide,  and  is  widely  used  as  an  abrasive.  The  ruby  is  pure  alu- 
minium oxide  tinted  by  a  trace  of  a  compound  of  chromium,  while 
the  sapphire  is  the  same  material  colored  with  aluminate  of  cobalt. 
Both  can  be  made  artificially  by  adding  a  little  of  the  oxide  of  chro- 
mium or  of  cobalt  when  aluminium  oxide  is  produced  by  Goldschmidt's 
method  (p.  684).  The  alumina  made  by  gently  heating  the  hydroxide 
interacts  easily  with  acids,  but  after  being  strongly  heated  it  re- 
sembles natural  alumina  in  being  very  slowly  affected  by  them.  Min- 
erals containing  insoluble  compounds  of  aluminium  are  attacked  when 
heated  strongly  with  potassium  bisulphate  (cf.  p.  560),  the  sulphate  of 
aluminium  being  formed. 

Aluminium  Sulphate  :  The  Alums.  —  The  sulphate  is  prepared 
by  treating  either  the  hydroxide  or  pure  clay  (kaolin)  with  sulphuric 
acid.  In  the  latter  case  the  insoluble  residue  of  silicic  acid  is  removed 
by  nitration : 

H2Al2(Si04)2  +  3H2S04  ->  A12(S04)8  +  2H2Si03  +  2H20. 

The  salt  crystallizes  from  water  as  A12(S04)3, 18H20,  forming  aggre- 
gates of  leaflets  which  are  very  soluble.  The  solution  is  acid  in 
reaction.  This  compound  is  used  as  a  mordant  under  the  name 
of  "concentrated  alum."  It  is  employed  also  in  sizing  cheaper 
grades  of  paper,  an  operation  required  to  prevent  the  absorption  and 
consequent  spreading  of  the  ink.  For  writing-paper,  gelatine  solution 
is  employed.  In  making  printing-papers,  rosin  soap  (made  by  dissolv- 
ing rosin  in  caustic  soda)  is  mixed  with  the  pulp,  and  aluminium 
sulphate  is  added.  The  rosin  and  aluminium  hydroxide  are  precipi- 
tated, perhaps  in  feeble  combination,  and  pressing  between  hot  rollers 
afterwards  melts  the  former  and  gives  a  surface  to  the  paper. 

When  sulphate  of  potassium  is  added  to  a  strong  solution  of  alu- 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS     687 

minium  sulphate,  octahedral  crystals  of  alum  (potash  alum,  K2S04,  A12 
(S04)3,24H2O)  are  deposited.  This  is  a  double  salt,  and  is  one  of  a 
large  number  known  as  the  alums.  These  have  the  general  formula 
M2IS04,M2II1(S04)3, 24H20,  and  may  be  made  as  above  by  using  a  sulphate 
of  a  univalent  metal  with  one  of  a  trivalent  metal.  Thus,  for  M1  we 
may  use  K,  NH,,  Eb,  Cs,  and  Tl1,  and  for  M111,  Al,  Fem,  Crm,  Mnm, 
and  Tlm.  We  may  even  employ  selenates,  such  as  K2Se04.  All  of 
the  resulting  double  salts  are  isomorphous,  and  a  crystal  of  one  will 
continue  to  grow  in  a  solution  of  another,  acquiring,  of  course,  an  outer 
layer  of  different  composition  but  of  the  same  crystallographic  orien- 
tation. 

Potassium-Aluminium,  Sulphate.  —  Ordinary  alum  K2S04, 
A12(S04)3,  24H20  is  made  from  aluminium  sulphate  obtained  from  clay 
(see  above).  It  is  also  prepared  by  heating  alunite,  a  basic  alum 
found  near  E-ome  and  in  Hungary,  and  extracting  the  product  with 
hot  water.  The  alunite,  having  the  composition  KA18(OH)6(S04)2, 
leaves  an  insoluble  residue  of  the  hydroxide,  mixed  with  ferric  oxide 
which  is  present  as  an  impurity : 

2KAl3(OH)6(S04)2 1>  K2S04,A12(S04)3  +  4A1(OH)3. 

The  aqueous  solution  of  alum  contains,  at  10°,  9  parts  of  the  anhy- 
drous salt  in  100  parts  of  water,  and  at  100°  422  parts  in  100  of  water. 
The  hydrated  salt  melts  at  90°.  An  aqueous  solution  of  this  salt  or  of 
sodium  phosphate  (p.  577)  is  used  for  fire-proofing  draperies,  because  the 
crystals  deposited  in  the  fabric  melt  easily,  and  the  fused  material 
protects  the  fibers  from  access  of  oxygen.  When  heated  more  strongly 
alum  loses  its  water  of  hydration  together  with  some  sulphur  trioxide, 
and  leaves  a  slightly  basic,  anhydrous  salt  known  as  "  burnt  alum." 
A  solution  of  alum  dissolves  a  considerable  amount  of  aluminium  hy- 
droxide, giving  "  neutral  alum,"  a  basic  salt  K2SO4,  A14(OH)6(S04)3  used 
as  a  mordant.  The  substance  is  usually  prepared  by  adding  sodium 
carbonate  to  the  solution  of  alum  as  long  as  the  aluminium  hydroxide, 
formed  locally,  continues  to  redissolve. 

Aluminium  Sulphide.  —  This  compound  is  most  easily  obtained 
by  mixing  pyrite  with  aluminium  powder  and  igniting  with  magnesium 
ribbon  (p.  684) : 

3FeS2  +  4A1  ->  2A13S8  +  3Fe. 


688  INORGANIC   CHEMISTRY 

It  forms  a  grayish-black  solid,  and  is  decomposed  by  water  like  mag- 
nesium sulphide,  giving  the  hydroxide  and  hydrogen  sulphide. 

Aluminium  Acetate.  —  This  salt  is  used  by  dyers,  because,  being 
a  salt  of  a  weak  base  and  a  weak  acid,  it  is  much  hydrolyzed  by  water, 
especially  at  100°.  In  mordanting,  it  thus  gives  aluminium  hydroxide 
very  easily.  It  is  made  by  treating  lead  or  barium  acetate  with 
aluminium  sulphate,  and  filtering  and  crystallizing  the  solution : 

A12(SO4)3  +  3Ba(C2H302)2  <=±3BaSOJ+2Al(C2H3O2)3. 

Dyeing :  Mordanting.  —  The  problem  of  the  dyer  is  to  confer 
the  desired  color  upon  a  fabric  made,  usually,  of  cotton,  linen,  wool, 
or  silk,  and  to  do  this  in  such  a  way  that  the  dye  is  fast  to  (i.e.,  is  not 
removed  or  destroyed  by)  rubbing,  and  often,  also,  to  washing  with 
soap.  To  understand  the  means  by  which  this  is  achieved,  it  must  be 
noted  that  cotton  and  linen  consist  of  hollow  fibers  of  the  composition 
of  cellulose  (C6H10O6),,..  Wool  is  made  of  hollow  fibers,  also,  and  silk  of 
rods,  but  the  material  is  entirely  different.  It  contains  17  per  cent 
of  nitrogen  in  the  case  of  wool,  and  20  per  cent  in  the  case  of  silk,  and 
the  nitrogen  compounds  of  which  the  material  is  composed  are  much 
more  active  chemically  than  is  cellulose,  and  combine  incomparably  more 
easily  and  firmly  with  the  many  kinds  of  organic  compounds  which 
are  used  as  dyes.  Hence,  stains  on  wool  and  silk  are  much  less  often 
removable  by  washing  than  are  those  on  cotton. 

We  have  space  to  mention  only  three  kinds  of  dyes  and  to  describe 
in  mere  outline  their  use.  These  are: 

1.  Insoluble  colored  bodies  which  are  formed  by  precipitation 
within  the  fibers  and  may  be  applied  to  any  fabric,  for  their  retention 
is  due  to  mechanical  and  not  to  chemical  causes.  If  cotton  is  boiled 
in  a  solution  of  lead  acetate  (or,  better  still,  sodium  plumbite,  q.v.\ 
and  is  then  soaked  in  boiling  potassium  chromate  solution,  it  is  dyed 
a  brilliant  and  permanent  yellow.  Lead  chromate  is  the  colored  body  : 

Pb(C2H802)2  +  K2Cr04  <±  2KC2H302  +  PbCr04  J. 

In  indigo  dyeing  the  fabric  is  saturated  with  a  solution  of  indigo- white 
(an  acid  substance)  in  caustic  soda,  and  is  then  exposed  to  the  air. 
Indigo-blue  is  formed  by  oxidation,  and,  being  insoluble,  is  precipitated 
in  the  fibers : 

2C16H12N202  +  02  -»  2C16H10N202  J  +  2H20. 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS     689 

2.  We  have  direct  or  substantive  dyes,  which  are  withdrawn  from 
a  solution  by  the  goods  which  are  being  dyed,  and  confer  upon  the  latter 
a  depth  of  color  depending  on  the  strength  of  the  solution  and  the  affinity 
of  the  material  for  the  dye.     The  union  is  due  in  some  cases  to  chemical 
combination,  and  in  others  to  the  fact  that  the  dye  is  more  soluble  in 
the  material  being  dyed  than  in  water,  and  gives  a  solid  solution  in  the 
former  (cf.  pp.  146,  235).     When  the  case  is  simply  one  of  extraction  of 
the  dye,  due  to  the  solvent  power  of  the  goods,  the  dye  must  necessarily  be 
removed  again  by  washing  with  sufficient  water.    Only  a  small  minority 
of  direct  dyes  are  taken  up  by  cotton  or  linen  in  such  a  way  that  they 
cannot  be  washed  out.     Congo  red,  C32H22N6S206Na2,  is  soluble  in  water, 
and  is  used  in  dyeing  both  cotton  and  wool.     The  dye  is  much  faster 
on  the  latter  than  on  the  former,  however. 

3.  The  last  class  comprises  the  mordant  or  adjective  dyes.     They 
work  on  the  principle  that  the  cloth  is  first  impregnated  with  a  substance 
capable  of  attaching  itself  both  to  the  cloth  and,  subsequently,  to  the 
dye  also,  and  is  then  immersed  in  the  dye  itself.     Substances  of  this 
kind  are  tannic  acid  (for  basic  dyes)  and  colloidal  hydroxides  (for  acid 
dyes)  like  those  of  aluminium,  tin,  iron,  and  chromium.     They  are 
called  mordants  (Lat.  mordere,  to  bite).     When  aluminium  hydroxide 
is  to  be  used,  the  cloth  is  first  treated  with  a  hot  solution  of  neutral 
alum,  aluminium  sulphate,  aluminium  acetate,  or  sodium  aluminate, 
and  thereby  acquires,  either  by  adsorption  or  feeble  combination,  a 
certain  amount  of  the  hydroxide.     The  fabric  is  then  boiled  in  water 
with  the  dye.     If,  for  example,  alizarin  (madder)  is  used,  the  cloth 
is   dyed   Turkey   red.      Alizarin   is   an   orange-yellow,  very   slightly 
soluble   acid  of    the    composition    C14H804.     Since  the    color  is  that 
of  the  compound  of  the  dye  with  the  mordant,   different   mordants 
give    different    colors,    or    shades    of    color,    with    the    same    dye. 
This    may    be    illustrated    by    preparing    three    solutions,    one    of 
ferric  chloride  (q.v.)  containing  ferric  hydroxide  in  solution,  one  of 
aluminium  acetate,   and  one  of  chromium   acetate.      When   a   drop 
or  two  of  an  alcoholic  solution  of  alizarin  is  added  to  each,  a  precipi- 
tate at  once  appears,  which  in  the  first  case  is  violet,  in  the  second 
bright-red,  and  in  the  third  claret-red.     These  insoluble  compounds  of 
dyes  with  mordants  are  identical  with  the  coloring  matters  produced 
in  the  cloth,  and  are  called  lakes  (Fr.  laque}  lac). 

Kaolin  and  Clay :    Earthenware  and  Porcelain.  —  By   the 

action  of  water  and  carbon  dioxide  upon  granite  and  other  rocks  con- 


690  INORGANIC   CHEMISTRY 

taining  feldspar  KAlSi8O8,  the  potash  is  slowly  removed,  and  the  com- 
pound changed  largely  into  a  hydrated  orthosilicate  H2Al2(Si04)2,H20. 
When  this  remains  in  situ,  it  forms  kaolin  or  china  clay,  a  white, 
crumbly  material.  It  usually  contains  particles  of  mica  and  free  silica. 
When  washed  away  and  redeposited,  it  acquires  compounds  of  iron,  and 
more  or  less  of  the  carbonates  of  calcium  and  magnesium,  becoming  com- 
mon clay.  Ocher,  umber,  and  sienna  are  clays  colored  with  oxides  of 
iron  and  manganese.  Fuller's  earth  is  a  purer  variety. 

On  account  of  its  plasticity  when  moist,  and  its  tendency  to  become 
hard,  but  not  to  melt,  when  heated  strongly,  clay  is  used  in  making 
bricks,  pottery,  and  porcelain.  The  presence  of  calcium  and  magne- 
sium carbonates  makes  the  clay  more  fusible,  that  of  silica  less  so.  Iron 
compounds  cause  it  to  turn  red  during  firing.  The  impure  varieties 
are  formed  into  bricks  and  tiles,  and  are  fired  at  a  low  temperature. 
The  efflorescence  which  often  appears  on  the  surface  of  the  bricks 
("niter")  is  generally  due  to  sodium  sulphate  or  sodium  chloride 
present  originally  in  the  clay.  For  earthenware,  glazing  must  be 
applied  to  make  the  vessels  water-tight.  This  is  often  done  by  throw- 
ing salt  into  the  kiln.  The  hot  steam  hydrolyzes  the  salt  to  sodium 
hydroxide  and  hydrochloric  acid,  and  the  former  combines  with  the 
clay,  giving  a  fusible  silicate  which  fills  the  pores  of  the  surface.  For 
porcelain,  very  pure  clay,  free  from  iron,  is  employed,  and  it  is  mixed 
with  feldspar  and  quartz.  The  feldspar  melts  and  fills  the  pores  so  that 
a  continuous,  semi-transparent  material  results.  For  china  painting, 
powdered  enamels  (p.  607)  and  metallic  oxides  which  combine  with 
the  clay,  giving  colored  silicates,  are  used. 

Porcelain,  if  made  with  sufficient  silica,  is  very  infusible.  It  is 
attacked  by  aqueous  and  by  fused  alkalies,  however,  giving  soluble  sili- 
cates. 

Ultramarine  was  formerly  obtained  by  pulverizing  natural  lapis 
lazuli.  Artificial  ultramarine  of  similar  composition  and  more  beauti- 
ful color  is  now  manufactured.  It  is  made  by  heating  together  kaolin, 
sodium  carbonate,  sulphur,  and  charcoal.  The  resulting  green  mass  is 
then  powdered,  mixed  with  sulphur,  and  heated  again  until  it  acquires 
the  desired  shade.  The  product  is  used  largely  in  making  wall-papers, 
water-color  paints,  and  laundry  blue.  Note-paper  is  often  colored 
with  it.  Its  composition  is  approximately  4NaAlSiO4,  Na^Sg,  but  the 
cause  of  the  brilliant  color  is  not  understood.  Hydrochloric  acid  lib- 
erates sulphur  and  hydrogen  sulphide  and  destroys  the  color  of  ultra* 
marine,  decomposing,  apparently,  the  sulphide  of  sodium. 


ALUMINIUM  AND  THE  METALS  OF  THE  EARTHS     691 

Analytical  Reactions  of  Aluminium  Compounds.  -  -The 
alkalies,  and  alkaline  solutions  like  that  of  ammonium  sulphide,  pre- 
cipitate the  white  hydroxide.  The  product  is  soluble  in  excess  of  the 
active  alkalies.  Soluble  carbonates  also  throw  down  the  hydroxide. 
Aluminium  compounds,  when  heated  strongly  in  the  flame  with  cobalt 
salts,  give  a  blue  aluminate  of  cobalt. 

Exercises.  —  1.  What  are  the  differences  between  zinc  and  alu- 
minium, and  their  corresponding  compounds  ? 

2.  Construct  equations  showing  (a)  the  hydrolysis  of  aluminium 
sulphate  (p.  686),  (b)  the  interaction  of  aluminium  sulphate  and  cobalt 
nitrate  in  the  Bunsen  flame. 

3.  Formulate  the  ionization  of  aluminium  hydroxide  (pp.  648,  684). 


CHAPTER   XL 
GERMANIUM,   TIN,   LEAD 

THE  elements  of  the  fifth  column  of  the  periodic  table,  aside  from 
carbon  and  silicon,  are  germanium  (Ge,  at.  wt.  72.5),  tin  (Sn,  at.  wt. 
119),  and  lead  (Pb,  at.  wt.  206.9).  These  are  on  the  right  side,  while 
titanium  (Ti,  at.  wt.  48.1),  zirconium  (Zr,  at.  wt.  90.6),  cerium  (Ce, 
at.  wt.  140.25),  and  thorium  (Th,  at.  wt.  232.5)  occupy  the  left  side. 

The  Chemical  Relations  of  the  Family.  —  All  of  these  ele- 
ments show  a  maximum  valence  of  four.  Germanium,  tin,  and  lead 
are  also  bivalent.  In  this  respect  they  resemble  carbon  and  differ 
from  silicon,  which  is  more  closely  allied  to  the  elements  on  the  left 
side  of  the  column.  The  oxides  and  hydroxides  in  which  these  three 
elements  are  bivalent  become  more  basic,  and  the  elements  themselves 
more  metallic  in  chemical  relations,  with  increase  in  atomic  weight. 
In  this  they  resemble  the  potassium,  calcium,  and  gallium  families. 
Curiously  enough,  the  same  three  hydroxides  are  also  acidic.  They 
are  more  strongly  acidic  than  is  zinc  hydroxide,  for  the  salts  they  form 
by  interaction  with  bases  are  less  hydrolyzed  than  are  the  zincates. 
This  acidic  character  likewise  increases  in  the  order  in  which  the 
elements  are  named  above. 

GERMANIUM. 

Germanium  (cf.  p.  412)  may  be  described  as  a  transition  element 
between  carbon  and  tin.  It  forms  two  oxides  GeO  and  Ge02  corre- 
sponding to  those  of  carbon  and  of  tin.  Germanious  oxide  is  not  very 
definitely  basic  or  acidic,  and  the  sulphide  is  the  only  other  well-de- 
fined compound  of  this  set.  Germanic  oxide  and  hydroxide  are  acidic 
entirely.  The  resemblance  to  carbon  is  shown  in  the  formation  of  an 
unstable  compound  with  hydrogen,  and  of  germanium  chloroform 
GeHCl8.  Like  carbon,  tin,  and  silicon,  germanium  gives  a  volatile 
chloride  GeCl4  (b.-p.  87°).  Like  tin  and  gold  (p.  638),  it  forms 
complex  sulphides  derived  from  germanic  sulphide,  such  as  K2GeS3. 
The  element  was  discovered  (in  1886)  in  argyrodite,  a  complex  sul- 
phide 4Ag2S,GeS2. 


GERMANIUM,    TIN,   LEAD  693 

TIN. 

The  Chemical  Relations  of  the  Element.  —  Tin  is  both  biva- 
lent and  quadrivalent.  Each  of  the  oxides  and  hydroxides  SnO  and 
Sn(OH)2,  Sn02  and  SnO(OH)2  (or  Sn(OH)4),  is  both  basic  and  acidic, 
so  that  there  are  really  four  series  of  compounds.  Still,  stannous 
hydroxide  is  mainly  a  base,  of  a  feeble  sort,  while  stannic  hydroxide  is 
mainly  an  acid.  jfThus  we  have  stannous  chloride,  sulphate,  and 
nitrate,  which  are  stable,  although  they  are  all  more  or  less  hydrolyzed 
by  water,  and  sodium  stannite  Na2.Sn02  which  is  unstable.  On  the 
other  hand,  stannic  nitrate,  sulphate,  and  chloride  are  completely 
hydrolyzed  by  water,  while  sodium  stannate  Na2SnO3  is  comparatively 
stable.  The  dioxide  SnO2  is  an  infusible  solid,  and  resembles,  there- 
fore, silicon  dioxide.  Tin  has  a  tendency  to  give  complex  acids  and 
salts,  like  H2SnCl6,  (NH4)2.SnCl6,  H2SnI6,  K2SnF6,  but  these  are 
ionized  also  to  a  small  extent  after  the  manner  of  double  salts,  giving 
ions  of  Sn'M*.  Tin  forms  no  compounds  with  hydrogen  and  no  salts 
with  weak  acids,  like  carbonic  acid. 

Occurrence  and  Extraction.  —  The  chief  ore  of  tin  is  tin-stone, 
or  cassiterite  Sn02,  which  consists  of  tetragonal  crystals  whose  dark 
color  is  due  to  the  presence  of  iron  compounds.  The  mineral  occurs 
in  Cornwall  and  the  East  Indies.  The  ore  is  roughly  pulverized 
and  washed,  to  remove  granite  or  slate  with  which  it  is  mixed,  and 
is  then  roasted,  to  oxidize  the  sulphides  of  iron  and  copper,  and  drive 
off  the  arsenic  which  it  contains.  After  renewed  washing  to  eliminate 
sulphate  of  copper  and  oxide  of  iron,  it  is  reduced  with  coal  in  a 
reverberatory  furnace.  The  tin  is  afterwards  remelted  at  a  gentle 
heat,  and  the  pure  metal  flows  away  from  compounds  of  iron  and 
arsenic.  In  1900  the  production  was  4100  tons  and  63,700  tons  in 
England  and  in  the  East  Indies,  respectively.  These  quantities 
together  constitute  83  per  cent  of  the  total  world's  output. 

Physical  and  Chemical  Properties.  —  Tin  is  a  silver-white, 
crystalline  metal  of  low  tenacity  but  great  malleability  (tinfoil).  Its 
specific  gravity  is  7.3,  and  its  melting-point  about  233°.  Tin  some- 
times changes  into  a  gray,  pulverulent,  specifically  lighter  modification 
(sp.  gr.  5.85)  when  it  is  kept  at  a  low  temperature.  The  transition 
point  is  20°  (cf.  Sulphur,  p.  368),  and  ordinary  tin,  although  it  can  be 
kept  almost  indefinitely,  is  therefore  really  in  a  inetastable  condition 
(p.  159)  below  this  temperature. 


594  INORGANIC   CHEMISTRY 

Tin-plate  (cf.  p.  673)  is  made  by  dipping  carefully  cleaned  sheets 
of  mild  steel  into  molten  tin.  Vessels  of  copper  are  also  coated, 
internally,  with  tin,  to  prevent  the  formation  of  the  basic  carbonate 
(p.  618).  For  this  purpose  they  are  cleaned  with  ammonium  chlo- 
ride, sprinkled  with  rosin  (to  reduce  the  oxide),  and  heated  to  230°. 
Molten  tin  is  then  spread  on  the  surface  with  a  piece  of  tow.  Com- 
mon pins  are  made  of  brass  wire,  and  are  coated  with  tin  by  being 
shaken  in  a  solution  containing  a  salt  of  this  metal.  The  zinc  in  the 
alloy  displaces  some  of  the  tin,  and  this  is  deposited  on  the  surface  of 
the  brass.  Alloys  of  tin,  such  as  bronze  (p.  619),  soft  solder  (50  per 
cent  lead),  pewter  (25  per  cent  lead),  and  britannia  metal  (10  per 
cent  antimony  and  some  copper),  are  much  used  in  the  arts. 

Tin,  although  it  displaces  hydrogen  from  dilute  acids,  is  not  tar- 
nished by  moist  air.  With  warm  hydrochloric  acid  it  gives  stannous 
chloride  SnCla  and  hydrogen.  Hot,  concentrated  sulphuric  acid  forms 
stannous  sulphate  SnS04  and  sulphur  dioxide.  Nitric  acid,  when  cold 
and  dilute,  interacts  with  it,  giving  stannous  nitrate,  and  a  portion  of 
the  nitric  acid  is  reduced  to  ammonia  (cf.  p.  446)  : 

4Sn  +  10HN08  ->  4Sn(N08)2  +  3H2O  +  NH4N08. 

With  concentrated  nitric  acid,  stannic  nitrate  is  formed,  but  most  of 
this  salt  is  hydrolyzed  by  the  water  at  the  high  temperature  of  the 
action  (cf.  p.  657),  and  metastannic  acid  (H2Sn08)6  remains.  The  final 
result  is  therefore  shown  by  the  equation  (simplified): 

Sn  +  4HN08  -4  H2Sn08  -f  4N02  +  H2O. 

The  white,  insoluble  product  continues  to  give  nitric  acid  during  pro- 
longed washing,  and  seems  therefore  to  contain  some  nitrate,  or  basic 
nitrate.  Tin  also  displaces  hydrogen  from  caustic  alkalies,  giving  a 
metastannate,  such  as  K2Sn08. 

Stannous  Chloride.  —  This  salt  is  made  by  the  interaction  of  tin 
and  hydrochloric  acid.  Evaporation  of  the  solution  gives  the  colorless 
SnCl2, 2H20.  When  the  crystals  are  heated,  or  when  a  strong  aqueous 
solution  is  diluted,  the  salt  is  partially  hydrolyzed.  In  the  latter  case 
the  basic  chloride  Sn(OH)Cl  is  deposited.  By  presence  of  excess  of 
hydrochloric  acid,  the  hydrolysis  is  prevented.  The  solution  is  used 
as  a  mordant  (p.  688). 


GERMANIUM,    TIN,    LEAD  695 

Stannous  chloride  tends  to  pass  into  stannic  chloride  SnCl4,  and  is 
therefore  an  active  reducing  agent.  Thus,  it  reduces  the  chlorides  of 
mercury  (p.  655)  and  of  the  noble  metals,  liberating  the  free  metals. 
The  action  is  of  the  form  Hg"  -f-  Sn"  — »Hg  +  Sn"".  Stannous 
chloride  reduces  cupric  and  ferric  chlorides  to  the  cuprous  and  ferrous 
conditions  in  like  manner  : 

2FeCl8  +  SnCL,  ->  2FeCL,  +  SnCl4   or    2Fe"'  +  Sn"  ->  2Fe"  +  Sn—'. 

It  also  reduces  free  oxygen,  or,  what  is  the  same  thing,  is  oxidized  by 
the  air.  In  this  case,  stannic  chloride  is  formed  in  the  acid  solution 
and  the  liquid  remains  clear ;  in  the  neutral  solution  a  precipitate  of 
the  basic  chloride  is  formed  as  well : 

6SnCl2  +  2H20  +  O2  ->  4Sn(OH)Cl  +  2SnCl4. 

Powdered  tin,  if  placed  in  the  bottle  along  with  the  acid  solution,  will 
undo  the  effects  of  this  action  by  reducing  the  stannic  salt  to  the 
stannous  condition  once  more. 

Stannic  Chloride.  —  When  chlorine  acts  upon  tin,  or  upon  stan- 
nous chloride  (either  solid  or  dissolved),  stannic  chloride  is  formed. 
The  compound  is  a  colorless  liquid  (b.-p.  114°)  which  fumes  very 
strongly  in  moist  air,  giving  hydrochloric  acid  and  stannic  acid.  It 
was  formerly  known,  after  its  discoverer  (1605),  as  spiritus  fumans 
Libavii.  The  aqueous  solution,  when  freshly  made,  has  almost  no 
conductivity,  and  the  compound  is  therefore  very  slightly  ionized. 
As  hydrolysis  proceeds,  the  conductivity  increases,  but  the  hydro- 
chloric acid  is  the  conducting  substance.  After  a  time  hydrolysis 
becomes  almost  complete.  The  stannic  acid  which  is  formed  is  not 
precipitated,  however,  but  remains  dissolved  in  colloidal  (p.  523) 
form: 

SnCl4  +  4H20  <±  4HC1  +  Sn(OH)4. 

The  chloride,  with  small  amounts  of  water,  gives  crystalline  hydrates 
SnCl4,3H20,  SnCl^HjjO,  and  SnCl^SHgO,  of  which  the  second  is  used 
as  a  mordant  under  the  name  "  oxymuriate  "  of  tin.  The  double  (or  per- 
haps complex)  salts  SnCl4,2MTCl  are  easily  made.  Ammonium-stan- 
nic chloride  or  "pink-salt"  (NH4)2SnCl6,  which  is  an  example  of  this 
class  of  compounds,  is  used  as  a  mordant  on  cotton.  It  gives  a  red 
lake  with  alizarine  (p.  689). 

Stannic  bromide  SnBr4  melts  at  30°,  boils  at  201°,  and  is  soluble  in 
water. 


696  INORGANIC   CHEMISTRY 

a-Stannic  Acid  and  its  Salts.  —  When  a  solution  of  stannic 
chloride  is  treated  with  ammonium  hydroxide,  a  white,  gelatinous  pre- 
cipitate is  formed.  To  this  the  formula  H2Sn08  is  generally  assigned  : 

SnCl4  +  4NH4OH  ->  4NH4C1  +  H2Sn08  +  H^O. 

It  is,  however,  in  reality,  a  hydrogele,  and  loses  water  gradually  until 
the  dioxide  remains.  Thus,  neither  Sn(OH)4  nor  SnO(OH)2  is  obtain- 
able as  a  definite  compound.  When  stannic  oxide  is  fused  with 
caustic  soda,  the  metastannate,  or  a-stannate,  is  formed  : 

Sn03  +  2NaOH  ->  Na2Sn03  +  H2O. 


This  compound  is  obtainable  as  Na^SnOg,  3H20,  and  is  used  as  a 
mordant  under  the  name  of  "  preparing  salt."  When  its  solution  is 
acidified,  the  above  mentioned  a-stannic  acid  is  formed  by  double  decom- 
position. This  a-stannic  acid  interacts  readily  with  acids  and  alkalies, 
and  the  chloride  obtained  from  it  is  identical  with  stannic  chloride 
described  above. 

The  a-stannates  of  the  metals,  aside  from  those  of  potassium  and 
sodium,  like  the  silicates  and  carbonates  which  they  much  resemble, 
are  all  insoluble  in  water,  and  may  be  made  by  double  decomposition. 

ft-Stannic  Acid,  or  Metastannic  Acid.  —  The  product  of  the 
action  of  nitric  acid  upon  tin  is  a  hydrated  stannic  oxide  like  the  fore- 
going substance,  but  is  not  identical  with  it.  It  is  not  easily  soluble 
in  alkalies.  By  boiling  it  with  caustic  soda,  however,  and  then  ex- 
tracting with  pure  water,  a  soluble  sodium  /8-stannate,  Na^SngOn,  is 
obtained,  /^-stannic  acid  is  also  very  slowly  attacked  by  acids,  and 
the  chloride  secured  from  it  is  not  identical  with  the  ordinary  chloride. 
For  these  reasons  it  is  supposed  to  be  a  hydrate  of  a  polymer  of 
stannic  oxide  (Sn02)B.  When  fused  with  caustic  soda,  it  gives  the 
same  a-stannate  as  does  the  dioxide  itself. 

The  difference  between  the  properties  of  the  two  stannic  acids 
was  noticed  by  Berzelius  (1811),  and  was  the  first  case  in  which  iden- 
tity in  composition  was  found  not  to  be  accompanied  by  identity  in 
properties  (cf.  Isomers,  p.  488). 

The  Oxides  of  Tin.  —  When  stannous  oxalate  is  heated  in  absence 
of  air,  stannous  oxide  remains  :  SnC204  —  >  SnO  4-  C02  +  CO.  It  is  a 
black  powder  which  burns  in  the  air,  giving  the  dioxide.  The  corre- 
sponding hydroxide  Sn20(OH)2  is  formed  by  adding  sodium  carbonate 


GEBMANIUM,   TIN,  LEAD  697 

to  stannous  chloride  solution.  It  is  a  white  powder,  easily  dehydrated, 
and  interacts  with  alkalies  to  give  a  soluble  stannite,  such  as  NaJSnO.,. 
When  the  solution  is  boiled,  tin  is  deposited,  and  sodium  stannate  is 
formed,  the  behavior  resembling  that  of  cuprous  oxide  when  heated 
with  acids  (p.  622).  With  acids,  the  hydroxide  gives  stannous  salts. 

Stannic  oxide  is  found  in  nature  (p.  693),  and  may  be  made  in  pure 
form  by  igniting  /8-stannic  acid.  When  heated,  it  becomes  yellow,  but 
recovers  its  whiteness  when  cooled  (cf.  Zinc  oxide,  p.  648).  Prepared 
at  a  low  temperature,  it  interacts  easily  with  acids,  but  after  strong 
ignition,  is  affected  by  them  very  slowly. 

The  Sulphides  of  Tin.  —  Stannous  sulphide  is  obtained  as  a  dark- 
brown  precipitate  when  hydrogen  sulphide  is  led  into  a  solution  of  a 
stannous  salt. 

Stannic  sulphide  is  formed  likewise  by  precipitation,  and  is  yellow 
in  color.  It  is  made  also  by  heating  together  tin  filings,  mercury,  sul- 
phur, and  ammonium  chloride.  The  mercury  and  ammonium  chloride 
are  ultimately  volatilized,  and  the  stannic  sulphide  remains  in  the  form 
of  yellow,  crystalline  scales  ("mosaic  gold"  or  " bronze  powder"). 
Stannic  sulphide  loses  sulphur  when  strongly  heated,  and  leaves  stan- 
nous sulphide.  It  is  not  much  affected  by  dilute  acids,  but  interacts 
with  solutions  of  ammonium  sulphide  or  sodium  sulphide,  giving  a 
soluble  complex  sulphide,  the  sulphostannate : 

SnS2  +  (NH4)2S  -»  (NH4)2.SnS8. 

The  corresponding  sodium  salt  is  easily  crystallized  in  the  form 
NajSnSg,  2H20.  Stannous  sulphide  is  not  affected  by  sulphides,  but 
polysulphides,  such  as  yellow  ammonium  sulphide,  give  with  it  the 
above  mentioned  sulphostannates : 

SnS  +  (NH4)2S  +  S  -»  (NH4)2.SnS3. 

With  acids  the  sulphostannates  undergo  double  decomposition,  but 
the  free  acid  H2.SnS8  thus  produced  is  unstable  and  breaks  up,  giving 
off  hydrogen  sulphide,  and  depositing  stannic  sulphide.* 

Analytical  Reactions  of  Salts  of  Tin.  —  The  two  ionic  forms 
of  tin,  Sn",  and  Sn"",  are  both  colorless.  Their  behavior  is  different. 

*  These  and  similar  compounds  are  often  called  thiostannates,  orthothian- 
timonates,  etc.  The  prefix  sulpho-  gives  more  euphonious  words,  however,  and 
is  used  here  for  all  excepting  the  thiocyanates. 


698  INORGANIC   CHEMISTRY 

They  give  a  brown  and  a  yellow  sulphide,  respectively,  with  hydrogen 
sulphide.  The  solubility  of  these  sulphides  in  yellow  ammonium 
sulphide  distinguishes  them  (cf.  p.  661)  from  those  of  cadmium,  copper, 
and  other  metals  whose  sulphides  are  similarly  insoluble  in  dilute  acids. 
The  sulphides  of  arsenic,  antimony,  and  gold  (q.v.  ),  however,  behave 
like  those  of  tin  in  this  respect.  The  reducing  power  of  distannion 
Sn**  is  very  characteristic  (p.  695).  Zinc  displaces  tin  from  solutions 
of  its  salts.  The  oxides  are  reduced  by  charcoal  in  the  reducing  part 
of  the  Bunsen  flame  and  the  metal  is  liberated. 

LEAD. 

The  Chemical  Relations  of  the  Element.  —  Lead  is  both  biva- 
lent and  quadrivalent.  The  oxides  PbO  and  Pb02,  and  the  correspond- 
ing hydrated  oxides,  are  both  basic  and  acidic.  Lead  monoxide  is  a 
fairly  active  base,  comparable  with  cupric  oxide,  and  lead  dioxide  a 
feeble  one.  Both  are  feebly  acidic.  The  salts  of  bivalent  lead,  like 
Pb(N03)2,  commonly  called  the  plumbic  salts,  are  somewhat  hydrolyzed 
by  water,  but  less  so  than  are  those  of  tin.  The  tetrachloride  and 
other  salts  of  quadrivalent  lead  are  completely  hydrolyzed.  The 
plumbites  Na^PbO-j  and  plumbates  Na2.Pb08  are  hydrolyzed  to  a  con- 
siderable extent.  All  the  compounds  in  which  lead  is  quadrivalent 
give  up  half  of  the  negative  radical  readily,  and  are  reduced  to  the 
"  plumbic  "  condition.  The  metal  displaces  hydrogen  with  difficulty, 
and  is  easily  displaced  by  zinc.  Lead  compounds  are  all  poisonous, 
and  the  effects  of  repeated,  very  minute  doses  are  cumulative  —  re- 
sulting in  "  lead  colic." 

Occurrence  and  Metallurgy.  —  Commercial  lead  is  almost  all 
obtained  from  galena  PbS,  which  crystallizes  in  cubes.  This  ore 
often  contains  considerable  amounts  of  silver  sulphide  Ag2S,  which 
is  isomorphous  with  it,  and  it  occurs  in  association  with  sulphides  of 
arsenic,  antimony,  zinc,  copper,  and  iron.  Other  salts  of  lead  are  of 
less  common  occurrence. 

The  sulphide  of  lead  is  first  roasted  until  a  sufficient  proportion 
of  it  has  been  converted  into  the  oxide  and  sulphate.  The  furnace- 
doors  are  then  closed,  and  the  temperature  raised  in  order  that  these 
products  may  interact  with  the  unchanged  part  of  the  sulphide  : 


PbS  +  2PbO 

PbS  +  PbS04  -+  2Pb  +  2S09. 


GERMANIUM,    TIN,    LEAD 

Another  plan  consists  in  heating  galenite  with  scrap  iron  or  iron  ores 
and  coal : 

PbS  +  Fe  ->  Pb  +  FeS. 

The  molten  ferrous  sulphide  rises  to  the  top  as  a  matte. 

The  purification  of  the  lead  from  the  other  metals  whose  sulphides 
have  been  reduced  at  the  same  time  is  often  troublesome.  In  Parke's 
process  (p.  627)  for  the  extraction  of  the  silver  by  means  of  zinc,  the 
greater  part  of  the  foreign  metals,  with  the  exception  of  bismuth, 
passes  into  the  zinc  scum.  About  0.5  per  cent  of  zinc  remains  in  the 
lead,  and  is  oxidized  by  the  action  of  a  jet  of  steam  before  the  lead  is 
poured  into  the  molds.  In  one  establishment,  at  Trail,  near  Eossland, 
B.C.,  the  refining  is  carried  out  electrolytically.  The  lead  is  cast  into 
plates,  and  the  process  is  similar  to  that  used  for  refining  copper 
(p.  617).  The  bath  is  a  solution  of  hydrofluosilicic  acid  containing 
hydrofluosilicate  of  lead.  The  less  electro-positive  metals,  Cu,  Sb,  Bi, 
As,  Ag,  Au,  with  10-16  per  cent  of  lead,  remain  as  a  sort  of  skeleton 
of  the  anode,  while  Zn,  Co,  Ni,  and  Fe  go  into  solution  and  are  not 
redeposited. 

Physical  and  Chemical  Properties.  —  Metallic  lead  is  gray  in 
color,  very  soft,  and  of  small  tensile  strength.  Its  specific  gravity  is 
11.4,  and  its  melting-point  326°.  While  warm,  it  is  formed  by  hydraulic 
pressure  into  pipes  which  are  used  in  plumbing  and  for  covering  elec- 
tric cables.  On  account  of  its  very  slow  interaction  with  most  sub- 
stances, sheet  lead  is  used  in  chemical  factories,  for  example,  to  line 
sulphuric-acid  chambers.  An  alloy  containing  0.5  per  cent  of  arsenic 
is  used  in  making  small  shot  and  shrapnel  bullets.  Type-metal  con- 
tains 20-25  per  cent  of  antimony  (q.v.).  In  both  cases  greater  hard- 
ness (cf.  p.  532)  is  secured  by  the  addition  of  the  foreign  metal. 

Lead  oxidizes  very  superficially  in  the  air.  The  suboxide  Pb2O  is 
supposed  to  be  first  formed.  The  final  covering  is  a  basic  carbonate. 
Contact  with  hard  waters  confers  upon  lead  a  similar  coating  composed 
of  the  carbonate  and  the  sulphate.  These  deposits,  being  insoluble, 
inclose  the  metal  and  protect  the  water  from  contamination  with 
lead  compounds.  Pure  rain-water,  however,  since  it  has  no  hard- 
ness, but  contains  oxygen  in  solution,  gives  the  hydroxide  Pb(OH)2, 
which  is  noticeably  soluble.  When  heated  in  the  air,  lead  gives 
the  monoxide  PbO  or  minium  Pb804,  according  to  the  temperature 
(see  below). 


700  INORGANIC   CHEMISTRY 

The  metal  displaces  hydrogen  from  hydrochloric  acid  very  slowly. 
It  is  hardly  affected  by  concentrated  sulphuric  acid,  although,  when 
the  commercial  acid  is  diluted  with  water,  a  slight  precipitate  of  lead 
sulphate,  acquired  from  the  evaporating-pans,  is  thrown  down.  Nitric 
acid  attacks  it  readily,  giving  lead  nitrate  and  oxides  of  nitrogen  (p. 
446). 

Chlorides  and  Iodide  of  Lead.  —  Plumbic  chloride  is  precipi- 
tated when  a  soluble  chloride  is  added  to  a  solution  of  a  soluble  lead  salt. 
It  is  slightly  soluble  in  water  (1.5  :  100)  at  18°,  and  considerably  more 
so  at  100°.  In  the  saturated  solution  at  25°  about  50  per  cent  of  the 
lead  is  in  the  form  Pb",  44  per  cent  as  PbCl*,  and  6  per  cent  as  PbCl2 
(of.  p.  346). 

Lead  tetrachloride  is  a  solid  at  —15°,  and  loses  chlorine  at  the 
ordinary  temperature.  It  is  made  by  passing  chlorine  into  plumbic 
chloride  suspended  in  hydrochloric  acid.  The  solution  appears  to  con- 
tain H2PbCl6.  With  ammonium  chloride  this  solution  deposits  crystals 
of  a  double  or  complex  salt  PbCl4,2NH4Cl  analogous  to  pink-salt 
(p.  695).  When  this  is  thrown  into  cold,  concentrated  sulphuric  acid, 
an  oil  of  the  composition  PbCl4  settles  to  the  bottom.  The  oil  fumes 
in  the  air,  and,  in  general,  closely  resembles  stannic  chloride  SnCl4. 
When  dissolved  in  little  water,  it  slowly  deposits  PbCl2  and  gives  off 
chlorine.  With  much  water  it  is  quickly  hydrolyzed,  and  lead  dioxide 
is  thrown  down : 

PbCl4  +  2H20  -»  Pb02  +  4HC1. 

Lead  iodide  PbI2  is  yellow  in  color,  and  is  formed  by  precipitation. 
It  is  somewhat  soluble  in  boiling  water,  and  crystallizes  in  yeHow 
scales  from  the  hot  solution. 

Plumbic  chloride  and  iodide  are  both,  under  some  conditions,  more 
soluble  in  acids  or  salts  with  a  common  negative  ion  than  they  are  in 
water,  and  form  soluble,  but  somewhat  unstable,  complex  salts  (<?/. 
p.  621). 

Oxides  and  Hydroxides.  —  There  are  five  different  oxides  of 
lead,  Pb20,  PbO,  Pb804,  Pb2O8,  and  Pb02.  The  suboxide  Pb20  is  a 
dark-gray  powder,  formed  by  gently  heating  the  oxalate.  Plumbic 
oxide,  or  lead  monoxide  PbO,  is  made  by  cupellation  (p.  627)  of  lead, 
and  the  solidified,  crystalline  mass  of  yellowish-red  color  is  sold  as 
"  litharge."  The  yellow,  powdery  form  is  called  "  massicot/7  and  may 


GERMANIUM,    TIN,    LEAD  701 

be  obtained  by  heating  the  nitrate  or  carbonate.  All  the  other  oxides 
yield  this  one  when  they  are  heated  above  600°  in  the  air.  Plumbic 
oxide  takes  up  carbon  dioxide  from  the  air,  and  therefore  usually  con- 
tains a  basic  carbonate.  It  dissolves  in  warm  sodium  hydroxide  solu- 
tion, giving  a  plumbite  Na,j.Pb02;  a  saturated  solution  redeposits 
part  of  the  oxide  in  crystalline  form  when  it  cools.  The  oxide  is  used 
in  glass-making  and  for  preparing  salts  of  lead. 

Plumbic  hydroxide  is  formed  by  precipitation.  It  gives  up  water 
in  three  stages  with  different  aqueous  tensions  (cf.  p.  122),  the  prod- 
ucts in  the  order  of  decreasing  tension  being  Pb(OH)2,  Pb2O(OH)2, 
Pb802(OH)2.  These  substances,  as  will  be  seen,  are  equivalent  in 
composition  to  PbO,  H2O,  2PbO,  H20,  and  3PbO,  H2O  respectively.  The 
hydroxide  is  observably  soluble  in  water,  and  gives  a  solution  with  a 
faintly  alkaline  reaction.  With  acids  it  forms  salts  of  lead.  It  inter- 
acts also  with  potassium  and  sodium  hydroxides  to  form  the  soluble 
plumbites,  like  Na^PbOg. 

Minium,  or  red  lead,  Pb304,  gives  off  oxygen  when  heated  : 


The  dissociation  pressure  varies  with  the  temperature  : 

Temperature     ........     445°         500°         555°         636° 

Pressure  in  mm  ........         5  60          183          763 

Since  the  partial  pressure  of  oxygen  in  the  air  is  150  mm.,  the  sub- 
stance decomposes  at  about  550°.  It  can  be  formed  in  air  by  reversal 
of  the  action  represented  above,  but  only  below  this  temperature 
(cf.  p.  591).  In  pure  oxygen  of  one  atmosphere  pressure  it  could  be 
formed  at  600°,  but  not  at  650°.  On  account  of  unequal  heating 
during  manufacture,  commercial  red  lead  is  never  fully  oxidized,  and 
always  contains  litharge.  Conversely,  commercial  litharge  usually 
contains  a  little  minium. 

Minium,  when  heated  with  warm,  dilute  nitric  acid,  is  decomposed, 
and  leaves  lead  dioxide  as  an  insoluble  powder.  It  is  therefore  re- 
garded as  lead  orthoplumbate  (see  below)  : 

Pb2.Pb04  +  4HN03  <±  2Pb(N08)2  +  H4Pb04. 

The  double  decomposition  as  a  salt  that  it  thus  undergoes  is  followed 
by  dehydration  of  the  plumbic  acid,  which  is  unstable  (H4PbO4  —  > 
Pb02  -f  2H20),  and  the  dioxide  remains.  Red  lead  is  used  in  glass- 


702  INORGANIC  CHEMISTRY 

making,  and,  when  mixed  with  oil,  gives  a  red  paint  which  is  specially 
applicable  to  iron-  work  (cf.  p.  658). 

Lead  dioxide  may  be  obtained  ^as  described  above  in  the  form  of  a 
brown  powder.  Unlike  most  oxides,  it  is  a  conductor  of  electricity. 
It  is  usually  made  by  adding  bleaching  powder  to  an  alkaline  solution 
of  plumbic  hydroxide  : 


Ca(OCl)Cl  +  H2O  ->  2STaOH  +  CaCL,  +  Pb02J  . 

In  this  action  we  may  regard  the  free  lead  hydroxide,  formed  by 
hydrolysis  of  the  plumbite,  as  being  oxidized  by  the  bleaching  powder. 
This  dioxide  is  an  active  oxidizing  agent.  It  interacts  with,  and 
sets  fire  to,  a  stream  of  hydrogen  sulphide,  and  it  liberates  chlorine  from 
hydrochloric  acid.  With  a_cids  J^  gives  no  hydrogen  peroxide,  and  is 
not  a  peroxide  in  the  restricted  sense  of  the  term  (p.  308).  Lead 
dioxide  interacts  with  potassium  and  sodium  hydroxides,  giving  soluble 
plumbates.  These  are  derived  from  metaplumbic  acid.  The  potas- 
sium salt  K2Pb08,3H20  is  analogous  to  the  metastannate  K2Sn03,3H2O 
(p.  696).  A  mixture  of  calcium  carbonate  and  lead  monoxide  absorbs 
oxygen  when  heated  in  a  stream  of  air,  and  the  yellowish-red  calcium 
orthopluinbate  is  formed  : 

4CaCOs  +  2PbO  +  O2  <±  2Ca2Pb04  +  4C02. 

The  action  is  reversible,  and  is  at  the  basis  of  Kassner's  'method  of 
manufacturing  oxygen  from  the  air. 

The  Storage  Battery.  —  In  the  storage  battery  the  plates  con- 
sist of  leaden  gratings,  the  openings  of  which  are  filled  with  the  active 
materials,  and  the  fluid  is  dilute  sulphuric  acid.  When  the  battery 
discharges,  the  SO/'  ions  migrate  towards  those  plates  (usually  the 
outer  ones)  which  are  filled  with  finely  divided  lead,  and  convert  this 
into  a  mass  of  (insoluble)  lead  sulphate  :  SO/'  +  Pb  —  >  PbSO4  4-  2  ©  . 
These  plates  receive  therefore  negative  charges.  Simultaneously  the 
H*  ions  pass  towards  the  other  plates,  and  reduce  to  monoxide  the 
lead  dioxide  with  which  they  are  filled  : 

2H%  -f  Pb02  ->  ELjO  +  PbO  +  2  ©. 

These  plates  acquire  positive  charges,  and,  by  secondary  interaction  of 
fche  monoxide  with  the  sulphuric  acid,  become  filled,  like  the  negative 


GERMANIUM,    TIN,   LEAD  703 

plates,  with  lead  sulphate.  During  the  discharge  the  fluid  thus  loses 
much  sulphuric  acid,  and  acquires  a  lower  specific  gravity,  a  fact  by 
means  of  which  the  approach  of  complete  Discharge  may  be  ascer- 
tained. The  E.M.F.  of  the  current  is  about  2  volts. 

When  a  current,  opposite  in  direction  to  the  one  which  it  yields,  is 
led  into  the  exhausted  cell,  the  negative  terminal  of  the  dynamo  cir- 
cuit being  connected  with  the  negative  pole  of  the  battery  and  the 
positive  with  the  positive,  a  new  set  of  changes  occurs.  The  H*  ions 
of  the  bath  are  attracted  to  the  plate  which  has  the  negative  charge,  and 
an  equivalent  number  of  SO/'  ions  are  formed,  so  that  only  metallic 
lead  remains  : 

PbS04  +  2H'  +  2  0  ->  Pb  +  2H-  +  SO/'. 

To  express  this  otherwise,  the  hydrogen  here  reduces  the  lead  sulphate 
(which  is  white)  to  metallic  lead  (a  black  powder).  Simultaneously 
the  SO/'  is  attracted  to  the  positively  charged  plate,  and  forms  lead 
persulphate  with  the  lead  sulphate  there  present :  SO/'  -f  PbSO4  -f-  2 
©  — »  Pb(S04)2.  The  lead  persulphate,  however,  in  a  battery  which  is 
working  normally,  is  at  once  hydrolyzed,  and  the  filling  of  the  plate  is 
changed  into  lead  dioxide  :  Pb(S04)2  -f  2H20  -+  Pb02  -f  2H2S04. 
Both  plates  are  thus  brought  back  to  the  condition  in  which  they  were 
before  the  discharge. 

The  set  of  changes  last  described,  that  involved  in  the  operation  of 
charging,  consumes  energy,  while  the  changes  connected  with  the  dis- 
charging, liberate  energy.  The  whole  may  be  put  into  a  single  equa- 
tion : 

Charge  — > 

2PbS04  +  2H2O  <=±  Pb  +  2H2S04  +  Pb02. 

< —  Discharge 

•  Lead  Nitrate.  —  This  salt  may  be  made  by  treating  lead,  lead 
monoxide,  or  lead  carbonate  with  nitric  acid.  It  forms  white,  anhy- 
drous octahedra.  The  nitrate  and  acetate  (see  below)  are  the  salts  of 
lead  which,  because  of  their  solubility,  are  most  commonly  used.  The 
solubility  of  the  nitrate  is,  48  parts  in  100  at  10°,  and  153  parts  in 
100  at  100°.  Since  the  solubility  increases  with  rise  in  temperature, 
we  should  expect  the  process  of  solution  to  be  accompanied  by  absorp- 
tion of  heat  (p.  260),  and  this  is  found  to  be  the  case.  On  account  of 
hydrolysis,  the  solution  is  acid  in  reaction. 


704  INORGANIC  CHEMISTRY 

Lead  Carbonate.  —  This  compound  is  found  in  nature  in  rhombic 
crystals  isomorphous  with  those  of  aragonite.  It  may  be  formed  as  a 
precipitate  by  adding  a  soluble  bicarbonate  to  lead  nitrate  solution. 
With  normal  sodium  carbonate,  a  basic  carbonate  Pb3(OH)2(C08)2  is  de- 
posited. This  basic  salt  is  identical  with  white  lead,  which,  on  account 
of  its  superior  opacity,  has  better  covering  power  than  zinc-white 
(p.  648)  or  permanent  white  (p.  610).  The  substance  is  manufactured 
in  various  ways,  all  of  which  involve  the  oxidation  of  the  lead  by  the 
air,  the  formation  of  a  basic  acetate  by  the  interaction  of  vinegar  or 
acetic  acid  with  the  oxide,  and  the  subsequent  decomposition  of  the  salt 
by  carbon  dioxide.  The  best  quality  is  obtained  by  the  Dutch  method. 
In  this,  gratings  of  cast  lead  are  placed  above  a  shallow  layer  of  vinegar 
in  small  pots.  These  pots  are  buried  in  manure,  which  by  its  decom- 
position furnishes  the  carbon  dioxide  and  the  necessary  warmth.  The 
gratings  are  gradually  converted  into  a  white  mass  of  the  basic  carbonate. 
The  vapor  of  acetic  acid  arising  from  the  vinegar  may  be  regarded  as  a 
catalytic  agent  (cf.  p.  508),  since  it  is  used  over  and  over  again. 

Lead  Acetate.  —  This  salt  is  made  by  the  action  of  acetic  acid  on 
litharge,  and  crystallizes  in  prisms  of  the  composition  Pb(C9H302)2, 
3H20.  It  is  easily  soluble  in  water,  and,  from  the  sweet  taste  of  the 
solution,  is  named  sugar  of  lead  (used  in  medicine).  The  basic  salt 
Pb(OH)(C2H302)  is  formed  by  boiling  a  solution  of  lead  acetate  with 
excess  of  litharge.  Unlike  most  basic  salts,  this  basic  salt  is  soluble  in 
water,  and  its  solution  has  a  faintly  alkaline  reaction. 

Lead  Sulphate.  —  The  sulphate  occurs  in  nature  as  anglesite,  and 
is  isomorphous  with  heavy  spar.  Being  insoluble  in  water,  it  is  easily 
obtained  by  precipitation.  It  is  slightly  soluble  in  concentrated  sul- 
phuric acid  (p.  387).  It  is  dissolved  to  a  noticeable  extent  by  nitric 
acid,  since  this  acid  is  more  active  than  sulphuric  acid  (cf.  p.  601). 
It  is  also  soluble  in  concentrated  sodium  hydroxide  solution,  on  ac- 
count of  the  removal  of  the  Pb"  ions  which  are  a  factor  in  its  solubility 
product  and  their  passage  into  the  Pb02"  anion  of  sodium  plumbite 
(cf.  p.  621).  Finally,  it  dissolves  easily  in  ammonium  tartrate,  since 
lead  enters  into  the  complex  anion  of  the  tartrates  in  the  same  way 
that  copper  does  (cf.  p.  623).  Barium  sulphate,  which  is  of  the  same 
order  of  insolubility  as  lead  sulphate,  is  somewhat  affected  by  nitric 
acid,  but  not  by  sodium  hydroxide  or  by  tartrates.  The  element 
barium  lacks  both  the  characteristics  which  lead  here  exhibits. 


GERMANIUM,   TIN,   LEAD  705 

Lead  Sulphide. —  Natural  lead  sulphide  (galena)  is  black,  and  its 
crystals  have  a  silvery  luster.  The  precipitated  salt  is  black  and 
amorphous.  It  is  more  easily  attacked  by  active  acids  than  is  mercuric 
sulphide  (cf.  p.  651).  Concentrated  nitric  acid,  being  an  oxidizing  agent 
as  well  as  an  acid,  attacks  and  dissolves  it  readily. 

Analytical  Reactions  of  Lead  Compounds.  —  Hydrogen  sul- 
phide precipitates  the  black  sulphide,  even  when  dilute  acids  are 
present.  Sulphuric  acid  throws  down  the  sulphate.  Potassium  hy- 
droxide gives  the  white  hydroxide,  which  dissolves  in  excess  to  form 
the  plumbite.  Potassium  chromate  or  dichromate  (q.v.)  gives  a  yellow 
precipitate  of  lead  chromate  PbCr04,  which  is  used  as  a  pigment  under 
the  name  of  "  chrome-yellow." 

TITANIUM,  ZIRCONIUM,  CERIUM,  THORIUM. 

The  metals  on  the  left  side  of  the  fifth  column  of  the  periodic  table 
are  all  quadrivalent,  although  compounds  in  which  a  lower  valence  ap- 
pears are  numerous  in  this  family.  The  first  two  are  feebly  base-form- 
ing as  well  as  feebly  acid-forming ;  the  last  two  are  base-forming  ex- 
clusively. 

Titanium  occurs  in  rutile  TiTi04.  Derived  from  it  are  a  number  of 
titanates  of  the  form  K2Ti03,  titanic  iron  ore  (menaccanite)  being  fer- 
rous titanate  FeTiO8. 

Zirconium  is  found  in  zircon,  the  orthosilicate  of  zirconium 
ZrSi04,  which  occurs  in  square  prismatic  crystals  isoniorphous  with 
rutile,  cassiterite  (SnSn04),  pyrolusite  (MnMnO4),  and  thorite 
(ThSi04).  The  oxide  is  used  in  making  the  incandescent  substance  in 
some  forms  of  gas  lamps. 

Cerium  occurs  chiefly  in  cerite  [Ce,  La,  Nd,  Pd]Si04,H20. 

Thorium  is  found  in  thorite  ThSi04  but  most  of  the  supply  comes 
from  monazite  sand.  The  nitrate  Th(N03)4, 6 H20  is  used  in  making 
Welsbach  incandescent  mantles  (cf.  Flame,  p.  510).  The  mantle  of 
knitted  cotton  is  dipped  in  a  solution  of  this  salt  along  with  one  per 
cent  of  cerium  nitrate  Ce(N03)4)  and  is  then  ignited.  The  oxides  ThO2 
(thoria)  and  Ce02  (ceria)  which  remain  form  a  fairly  coherent  mass. 
Larger  or  smaller  proportions  of  cerium  oxide  diminish  the  luminosity  of 
the  glowing  material.  It  is  supposed  that  the  particles  of  finely  divided 
cerium  oxide  assist  the  union  of  the  oxygen  and  illuminating-gas  cata- 
lytically  and  hence  are  themselves  raised  to  a  high  temperature  by  the 


706  INORGANIC   CHEMISTRY 

action  which  is  thus  concentrated  in  their  neighborhood.     The  thoria 
acts  simply  as  a  nonconducting  support  for  the  ceria. 

The  Nernst  lamp  is  an  incandescent  electric  lighting  arrangement  in 
which  a  rod  of  the  oxides  of  several  of  the  rare  metals  takes  the  place  of 
the  common  carbon  filament.  The  peculiarity  of  this  lamp  is  that  pre- 
heating is  required  before  the  rod  attains  a  temperature  at  which  it 
will  conduct  the  current.  When  this  point  has  been  once  reached,  the 
resistance  enables  the  current  to  maintain  the  rod  at  the  temperature 
of  incandescence.  For  equal  consumptions  of  electricity,  this  form  of 
electric  lamp  gives  a  greater  yield  of  light  than  does  the  ordinary, 
carbon-filament,  incandescent  bulb. 

Exercises.  —  1.  In  what  order  should  you  place  the  elements  dealt 
with  in  this  chapter,  beginning  with  the  least  metallic,  and  ending 
with  the  most  metallic  (p.  533)  ? 

2.  Construct  equations  showing  (a)  the  interaction  of  tin  and  con- 
centrated sulphuric  acid,  (b)  of   water   and   stannous   chloride,  (c)  of 
chlorauric  acid  and  stannous  chloride  (p.  695),  (d)  of  oxygen  and  stan- 
nous chloride  in  acid  solution,  (e)  the  decomposition  of  lead   oxalate 
(p.  700),  (/)  the  interaction  of  lead  monoxide  and  acetic  acid,  (g)  and 
of  lead  monoxide  and  lead  acetate. 

3.  To  which  class  of  ionic  actions  do  the  reductions  by  stannous 
chloride  belong  ? 

4.  What  interactions  probably  occur  when  lead  dioxide  liberates 
chlorine  from  hydrochloric  acid  ? 

5.  How  should  you  set   about   preparing  (a)  lead   oxalate  (insol- 
uble), (b)  lead  chlorate  (soluble)  ? 

6.  Should  the  formula   of  the  sulphate  of  quadrivalent  lead  be 
written  Pb(S04)2  or- PbS208,  and   is   it  related  to   persulphuric   acid 

(H2S208)  ? 

7.  Describe  in  terms  of  the  categories  used  in  connection  with  the 
phase  rule  (p.  592)  the  system  furnished  by  minium  at  500°. 


CHAPTER   XLI 
ARSENIC,   ANTIMONY,   BISMUTH 

THIS  family  is  very  closely  related  to  the  elements  phosphorus  and 
nitrogen  which  precede  it  in  the  same  column  of  the  periodic  table. 
In  reading  this  chapter,  therefore,  constant  reference  should  be  made 
to  the  chemistry  of  the  corresponding  compounds  of  phosphorus.  A 
general  comparison  of  the  elements  arsenic  (As,  at.  wt.  75),  antimony 
(Sb,  at.  wt.  120.2)  and  bismuth  (Bi,  at.  wt.  208.0)  with  each  other  and 
with  the  two  already  disposed  of  will  be  given  at  the  end  of  this  chap- 
ter. It  is  sufficient  here  to  say  that  arsenic  is  mainly  an  acid-forming 
element,  and  is  therefore  a  non-metal,  while  antimony  is  both  acid- 
forming  and  base-forming,  and  bismuth  is  base-forming.  Each  of  the 
three  elements  gives  sets  of  compounds  in  which  it  is  trivalent,  and 
others  in  which  it  is  quinquivalent.  None  of  the  free  elements  dis- 
places hydrogen  from  dilute  acids. 

ARSENIC. 

The  Chemical  Relations  of  the  Element.  —  Arsenic  forms  a 
compound  with  hydrogen  AsH3.  It  gives  several  halogen  derivatives  of 
the  type  AsX3  which  are  completely  hydrolyzed  by  water.  Its  oxides 
and  hydroxides  are  acidic.  Salts  derived  from  H3As03  (=  As208,3H20) 
and  HAs02  ( =  As203,H2O),  and  named  orthoarsenites  and  metarsenites, 
are  known,  as  are  also  salts  derived  from  various  arsenic  acids : 

H3As04  =  As205,3H20  =  AsO(OH)3, 
H4As207  =  As205,2H20  =  As203(OH)4, 
HAs03    =As205,H20    =AsO,(OH). 

In  analogy  to  the  phosphates  (p.  465),  these  are  named  ortho-,  pyro- 
and  metarsenates  respectively.  The  last  column  shows  a  method  of 
writing  the  formulae  which  is  often  adopted  (cf.  p.  470).  Sulphates, 
nitrates,  carbonates,  and  other  salts  of  arsenic  are  not  formed.  There 
are,  however,  complex  sulphides  of  the  forms  Na^.AsSg  (sodium  ortho- 
sulpharsenite),  Na8.AsS4  (sodium  orthosulpharsenate),  etc. 

In  many  natural  sulphides,  such  as  pyrite  FeS2  and  zinc-blende 

707 


708  INORGANIC    CHEMISTRY 

ZnS,  a  part  of  the  sulphur  is  replaced  by  arsenic,  which  must  there- 
fore be  playing  the  part  of  a  bivalent  element  in  such  cases.  When 
much  arsenic  is  present,  the  formulae  are  written  thus :  Fe[S,As]2  and 
Zn[S,As]. 

Occurrence  and  Preparation.  —  Arsenic  is  found  free  in 
nature.  It  occurs  also  in  combination  with  many  metals,  particularly 
in  arsenical  pyrites  FeAsS.  Two  sulphides  of  arsenic,  orpiment 
As2S3  and  realgar  As2S2,  and  an  oxide,  white  arsenic  As203,  are  less 
common. 

The  element  is  obtained  either  from  the  native  material  or  by  heating 
arsenical  pyrites  :  FeAsS  — >  FeS  •+•  As.  During  the  roasting  of  the 
sulphur  ores  of  metals,  arsenic  trioxide  is  formed  by  the  oxidation  of 
the  arsenic  so  frequently  present,  and  collects  as  a  dust  in  the  flues. 

Physical  Properties. —  In  its  ordinary  condition,  the  free  element 
is  steel-gray  in  color,  metallic  in  appearance,  and  crystalline  in  form. 
When  the  vapor  is  suddenly ,  cooled,  however,  a  yellow  variety  is 
obtained,  which  is  soluble  in  carbon  disulphide,  is  phosphorescent  in 
the  air,  and  in  other  ways  resembles  common  phosphorus.  This,  like 
yellow  phosphorus,  is  the  less  stable  form. 

Elementary  arsenic  is  easily  volatilized  at  180°,  and  acquires  a  vapor 
pressure  of  760  mm.  long  before  the  melting-point  (480°,  under  high 
pressure)  is  reached.  The  density  of  the  vapor  measured  at  644° 
gives  308.4  as  the  weight  of  the  G.M.V.  (22.4  liters  at  0°  and  760 
mm.).  The  weight  combining  with  one  unit  (35.45  g.)  of  chlorine,  is 
25  g.,  and  three  times  this  amount  of  the  element,  or  75  g.,  is  found  in 
the  G.M.V.  of  the  vapor  of  the  chloride.  It  is  also  the  smallest  weight 
found  in  the  G.M.V.  of  any  volatile  compound  of  arsenic,  and  is  there- 
fore accepted  as  the  atomic  weight.  Since  308.4  is  equal  approxi- 
mately to  4  x  75  (=  300),  the  formula  of  the  vapor  of  the  simple 
substance  at  644°  is  As4.  At  1700°  dissociation  has  occurred,  and  the 
formula  is  As2. 

Chemical  Properties.  —  The  free  element  burns  in  the  air,  pro- 
ducing clouds  of  the  solid  trioxide  As203.  It  unites  directly  with  the 
halogens,  with  sulphur,  and  with  many  of  the  metals.  When  boiled 
with  nitric  acid,  chlorine  water,  and  other  powerful  oxidizing  agents, 
it  is  oxidized  in  the  same  way  as  is  phosphorus,  and  yields  arsenic 
acid  H3As04. 


ARSENIC,    ANTIMONY,   BISMUTH  709 

Arsine.  —  This  substance  corresponds  in  composition  to  ammonia 
and  phosphine,  and  some  of  the  ways  in  which  it  may  be  formed  are 
analogous  to  those  used  in  the  case  of  these  substances.  Thus,  when 
arsenic  and  zinc  are  melted  together  in  the  proportions  to  form  zinc 
arsenide  Zn8As2,  and  the  product  is  treated  with  dilute  hydrochloric 
acid,  the  result  is  similar  to  the  action  of  water  or  dilute  acids  upon 
calcium  phosphide.  Arsine  in  fairly  pure  condition  is  evolved  as 

a  gas: 

Zn3As2  -f  6HC1  -4  2AsH3  +  3ZnCl2. 

Arsine  is  formed  also  by  the  action  of  nascent  hydrogen  (cf.  p.  423) 
upon  soluble  compounds  of  arsenic,  such  as  arsenious  chloride  AsCl8 
or  arsenic  acid.  When  a  solution  of  one  of  these  substances  is  added 
to  zinc  and  hydrochloric  acid  in  a  generating  flask,  the  disagreeable 
odor  of  the  hydrogen  evolved  shows  the  presence  of  the  arsine : 

AsCl8  +  3H2-»  AsH8  -|-  3HC1. 

This  method,  naturally,  does  not  furnish  pure  arsine,  for  free  hydrogen 
predominates  in  the  gas.  Pure  arsine  may  be  secured  by  leading  the 
mixture  with  hydrogen  through  a  U-tube  immersed  in  liquid  air.  The 
arsine  (b.-p.  —  40°)  condenses  as  a  colorless  liquid. 

Arsine  burns  with  a  bluish  flame,  producing  water  and  clouds  of 
arsenic  trioxide :  2AsH3  +  302  — >  3H20  -f-  As208.  The  combustion  of 
hydrogen  containing  arsine  produces  the  same  substances.  Since 
arsine,  when  heated,  is  readily  dissociated  into  its  constituents 
(cf.  p.  252),  the  vapor  of  free  arsenic  is  present  in  the  interior  of  the 
hydrogen  flame.  This  arsenic  may  be  condensed  in  the  form  of  a 
metallic-looking,  brownish  stain  by  interposition  of  a  cold  vessel  of  white 
porcelain.  Even  when  only  a  trace  of  the  compound  of  arsenic  has 
been  added  to  the  materials  in  the  generator,  the  stain  which  is  pro- 
duced is  very  conspicuous.  This  behavior  thus  furnishes  us  with  an 
exceedingly  delicate  test  —  Marsh's  test  —  for  the  presence  of  arsenic 
in  any  soluble  form  of  combination.  The  compounds  of  antimony 
alone  show  a  similar  phenomenon  (see  Stibine).  In  carrying  out  the 
test,  a  tube  of  hard  glass  is  attached  to  the  generator,  and  is  heated,  by 
means  of  a  Bunsen  flame,  at  a  point  near  to  the  flask.  With  this 
arrangement  the  arsenic  is  deposited  in  the  form  of  a  dark,  lustrous 
ring  just  beyond  the  heated  part.  Zinc  of  special  purity  must  be 
employed  for  generating  the  hydrogen,  as  all  common  specimens  of 
the  metal  contain  a  sufficient  amount  of  arsenic  to  give  the  metallic 


710  INORGANIC   CHEMISTRY 

film  without  any  special  addition  of  an  arsenic  compound,  and  a  blank 
experiment  must  be  run,  with  other  portions  of  the  same  reagents,  to 
guard  against  the  possibility  of  its  coming  from  any  of  them. 

Arsine  is  exceedingly  poisonous,  the  breathing  of  small  amounts 
producing  fatal  effects.  It  differs  from  ammonia  more  markedly  than 
does  phosphine,  for  it  is  not  only  without  action  on  water  and  on  acids, 
but  does  not  unite  directly  even  with  the  halides  of  hydrogen. 

Halides  of  Arsenic.  —  The  representatives  of  this  class  include  a 
liquid  trifluoride  AsF8,  a  pentafluoride,  which  is  obtained  only  as 
a  double  compound  with  potassium  fluoride,  a  liquid  trichloride  AsCl8, 
a  solid  tribromide  AsBr3,  and  a  solid  tri-iodide  AsI8. 

The  trichloride,  which  is  prepared  by  passing  chlorine  gas  into  a 
vessel  containing  arsenic,  is  easily  formed  as  the  result  of  a  vigorous 
action.  It  is  a  colorless  liquid,  boiling  at  130°*  When  mixed  with 
water  it  is  at  once  converted  into  the  white,  almost  insoluble  trioxide. 
The  action  is  presumably  similar  to  that  of  water  upon  the  correspond- 
ing compound  of  phosphorus  (p.  181),  but  the  arsenious  acid  for  the 
most  part  loses  water  and  forms  the  insoluble  anhydride : 

AsCl3  +  3H20  <=>  As(OH)8  +  3HC1, 
2As(OH)3  <=±  As203|  +  3H20. 

This  action,  however,  differs  markedly  from  the  other  in  that  it  is 
reversible,  and  arsenic  trioxide  interacts  with  aqueous  hydrochloric 
acid,  giving  a  solution  of  arsenious  chloride.  When  this  solution  is 
boiled,  the  volatility  of  the  arsenious  chloride  causes  it  to  be  carried 
over  with  the  hydrochloric  acid  (b.-p.  110°,  cf.  p.  182),  and  this  method 
of  separating  arsenic  from  other  substances  is  used  in  chemical 
analysis. 

O&ides  of  Arsenic;  —  Two  oxides  are  known  —  the  trioxide 
As203  and  the  pentoxide  As2O6.  Arsenic  trioxide  is  produced  by 
burning  arsenic  in  the  air  and  during  the  roasting  of  arsenical  ores 
(p.  708),  and  is  known  as  "  white  arsenic  "  or  simply  "  arsenic."  It  is 
purified  for  commercial  purposes  by  subliming  the  flue-dust  in  cylin- 
drical pots.  The  pure  trioxide  is  deposited  in  the  glassy  form  in 
the  upper  part  of  the  vessel.  It  passes  slowly  from  this  amorphous 
condition  into  the  common  crystalline  variety.  Its  vapor  density  in- 
dicates that  it  possesses  the  molecular  weight  As406,  but  the  simpler 
formula  expresses  its  chemical  properties  sufficiently  well. 


ARSENIC,    ANTIMONY,    BISMUTH  711 

When  treated  with  water,  the  trioxide  dissolves  to  a  very  slight 
extent  (0.3  :  100),  forming  arsenious  acid,  by  reversal  of  the  second  of 
the  equations  given  above.  As  nsual,  the  less  stable,  amorphous  variety 
is  the  more  soluble  (1  :  100).  In  boiling  water  the  solubility  is  eventu- 
ally much  greater  (11.5  :  100),  but  a  condition  of  equilibrium  is  reached 
very  slowly.  With  concentrated  sulphuric  acid  the  trioxide  forms  a  sul- 
phate of  rather  complex  composition,  indicating  that  it  has  basic  proper- 
ties, but  this  sulphate  is  decomposed  into  the  oxide  and  sulphuric  acid 
when  treated  with  water.  When  heated  in  a  tube  with  carbon,  this  ox- 
ide is  reduced,  and  the  free  element,  being  volatile,  is  deposited  upon 
the  cold  part  of  the  tube  just  above  the  flame.  It  is  an  active  poison, 
since  it  gradually  passes  into  solution,  forming  arsenious  acid. 

The  pentoxide  is  a  white  crystalline  substance,  formed  by  heating 
arsenic  acid  : 

2H3As04  ->  As205  +  3H20. 

When  raised  to  a  higher  temperature,  it  loses  a  part  of  its  oxygen, 
leaving  the  trioxide.  In  consequence  of  this  instability,  it  cannot  be 
formed  by  direct  union  of  oxygen  with  the  trioxide,  after  the  manner 
of  phosphorus  pentoxide. 

Acids  of  Arsenic.  —  When  elementary  arsenic  or  arsenious  ox- 
ide is  treated  with  concentrated  nitric  acid,  or  with  chlorine  and 
water,  arsenic  acid  is  produced.  The  substance,  which  is  a  deliques- 
cent white  solid,  possesses  a  composition  similar  to  that  of  orthophos- 
phoric  acid,  and  when  heated  loses  water  in  progressive  stages,  fur- 
nishing intermediate  acids  —  pyroarsenic  acid  and  metarsenic  acid  — 
and  finally  the  pentoxide.  The  relationship  of  the  substances 
(cf.  p.  465)  is  shown  by  the  formulae  :  H3As04  ->  H4As2O7  -»  HAs03 
— » As2O6.  These  acids,  however,  differ  from  the  corresponding  com- 
pounds of  phosphorus  in  that  upon  solution  in  water  they  immediately 
pass  back  into  the  ortho-acid.  The  final  elimination  of  all  the  water 
by  simple  heating  is  also  impossible  with  metaphosphoric  acid.  Many 
salts  of  these  acids  are  known,  those  of  orthoarsenic  acid  being  iso- 
morphous  with  the  corresponding  salts  of  phosphoric  acid.  The  red- 
dish-brown silver  orthoarsenate  and  the  white  MgNH4As04  resemble 
the  corresponding  phosphates  in  being  insoluble  in  water. 

Arsenious  acid,  like  sulphurous  and  carbonic  acids,  loses  water, 
and  yields  the  anhydride,  arsenic  trioxide,  when  the  attempt  is  made 
to  obtain  it  from  the  aqueous  solution.  The  potassium  and  sodium 


712  INORGANIC  CHEMISTRY 

arsenites,  K3As03  and  N"a3As03,  are  made  by  treating  arsenic  trioxide 
with  caustic  alkalies,  and  are  much  hydrolyzed  by  water.  The  arsen- 
ites of  the  heavy  metals  are  insoluble,  and  can  be  made  by  precipita- 
tion. Paris  green  and  Scheele's  green  (p.  624)  are  arsenites  of 
copper.  The  poisonous  effects  of  wall-paper  colored  with  these  com- 
pounds seem  to  be  due  to  volatile  organic  derivatives  of  arsine  which 
are  formed  by  the  action  of  a  mold.  In  cases  of  poisoning  by  white 
arsenic,  freshly  precipitated  ferric  hydroxide  or  magnesium  hydroxide 
is  administered,  since  by  interaction  with  the  arsenious  acid  they  form 
insoluble  arsenites.  The  salts  of  arsenious  acid  are  readily  oxidized, 
passing  into  arsenates.  The  action  of  a  standard  solution  of  iodine 
upon  sodium  arsenite,  for  example,  is  used  in  volumetric  analysis : 

Na3AsO3  +  H20  +  I2  <=±  Na3As04  +  2HI. 

Sulphides  of  Arsenic.  —  Three  sulphides  of  arsenic  are  known, 
As2S5,  As2S8,  As2S2.  The  first,  arsenic  pentasulphide,  is  obtained  as  a 
yellow  powder  by  decomposition  of  the  sulpharsenates  (see  below),  and 
by  leading  hydrogen  sulphide  into  a  solution  of  arsenic  acid  in  concen- 
trated hydrochloric  acid.  The  latter  action  seems  to  show  that  the 
ion  As""*,  derived  from  AsCl5,  is  present  in  the  solution. 

Arsenioua  sulphide  As2S8  occurs  in  nature  as  orpiment,  and  was 
formerly  used  as  a  yellow  pigment  (auripigmentum).  The  word 
arsenic  is  derived  from  the  Greek  name  for  this  mineral  (d/ao-en/cov). 
It  is  obtained  as  a  citron-yellow  precipitate  when  hydrogen  sulphide 
is  led  into  an  aqueous  solution  of  arsenious  chloride.  The  slowness 
with  which  the  precipitate  appears  is  due  to  the  temporary  formation 
of  a  colloidal  solution  of  the  sulphide.  When  arsenious  acid  solution 
is  employed,  the  precipitation  may  be  delayed  for  days.  A  beam  of 
light  falling  upon  the  solution  is  dispersed,  and  the  liquid  is  thus  shown 
to  be  filled  with  minute  particles,  which  ultimately  collect  in  flocculent 
form. 

Realgar  As2S2  is  a  natural  sulphide  of  orange-red  color,  and  is  also 
manufactured  by  subliming  a  mixture  of  arsenical  pyrites  and  pyrite : 

2FeAsS  +  2FeS2  -»  4FeS  +  As2S2f . 

It  burns  in  oxygen,  forming  arsenious  oxide  and  sulphur  dioxide,  and 
is  mixed  with  potassium  nitrate  and  sulphur  to  make  "  Bengal  lights." 

Sulpharsenites  and  Sulpharsenates.  —  The  sulphides  of  ar- 
senic interact  with  solutions  of  alkali  sulphides  after  the  manner  of 


ARSENIC,    ANTIMONY,    BISMUTH  713 

the  sulphides  of  tin  (p.  697),  giving  soluble,  complex  sulphides.  Arsen- 
ious  sulphide  with  colorless  ammonium  sulphide  gives  ammonium  sul- 
pharsenite,  and  with  the  yellow  sulphide  gives  ammonium  sulphar- 

senate : 

3(NH4)2S  +  As2S3  ->  2(NH4)8AsS3, 

3(NH4)2S  +  As2S3  +  2S  ->  2(NH4)8AsS4. 

Proustite  (p.  626)  is  a  natural  sulpharsenite  of  silver. 

There  are  also  metasulpharsenites  and  pyro-  and  metasulphar- 
senates  of  the  forms  KAsS2,  K4As2S7,  and  KAsSg.  In  fact,  there  are 
sulpho-compounds  corresponding  to  all  the  forms  of  the  oxygen  acids. 
The  particular  products  present  in  any  solution  depend  on  the  concen- 
trations used  in  making  it.  The  formation  of  these  soluble  compounds 
is  used  in  analysis  (cf.  p.  661). 

Since  these  salts  in  solution  furnish  the  ions  AsS3'"  and  AsS/", 
which  with  hydrogen  ions  give  a  feebly  ionized  sulpho-acid,  they  are 
decomposed  by  acids,  and  give  free  sulpharsenious  or  sulpharsenic 
acid  : 

(NH4)8AsS3  +  3HC1  -»  H3AsS3  +  3NH4C1, 
(NH4)3AsS4  +  3HG1  ->  H3AsS4  -f  3NH4C1. 

These  sulpho-acids,  however,  are  unstable,  and  at  once  break  up,  giving 
hydrogen  sulphide  as  a  gas,  and  the  sulphides  of  arsenic  as  yellow 
precipitates : 

2H3AsS3  ->  3H2S  +  As2S8, 

2H8AsS4  ->  3H2S  +  As2S5. 

ANTIMONY. 

The  Chemical  Relations  of  the  Element.  —  Antimony  resembles 
arsenic  in  forming  a  hydride  SbH3  and  halides  of  the  forms  SbX3  and 
SbX5.  The  latter  are  partially  hydrolyzed  by  water  with  ease,  but 
complete  hydrolysis  is  difficult  to  accomplish  with  cold  water.  The 
oxide  Sb208  is  basic  as  well  as  feebly  acidic,  and  the  oxide  Sb205  is 
acidic.  Aside  from  the  salts  derived  from  the  oxide  Sb203,  such  as 
Sb2(S04)8,  the  compositions  of  the  compounds  are  similar  to  those  of 
the  compounds  of  arsenic.  The  element  gives  complex  sulphides  like 
those  of  arsenic. 

Occurrence  and  Preparation.  —  Antimony  is  found  free  in 
nature  to  a  small  extent.  The  chief  supply,  however,  is  furnished 


714  INORGANIC   CHEMISTRY 

by  the  black  trisulphide  Sb2S3,  known  as  stibnite,  which  is  found  in 
Hungary  and  Japan,  and  forms  shining,  prismatic  crystals  of  the 
rhombic  system. 

Native  stibnite  is  roasted  in  the  air  in  order  to  remove  the  sulphur, 
and  the  white  oxide  which  remains  is  mixed  with  carbon  and  reduced 
by  strong  heat : 

Sb2S3  +  502  ->  Sb204  +  3S02, 

Sb2O4  +  40  ->  2Sb  +  400. 

Properties.  —  Antimony  is  a  white,  crystalline  metal.  It  is  brit- 
tle, and  easily  powdered.  Its  vapor  at  1640°  has  a  density  correspond- 
ing to  the  formula  Sb2,  while  at  lower  temperatures  Sb4  is  present. 
It  is  used  in  making  alloys  such  as  type-metal,  stereotype-metal,  and 
britannia  metal  (q.v.).  The  alloys  of  antimony  expand  during  solidifi- 
cation, and  therefore  give  exceptionally  sharp  castings. 

The  element  unites  directly  with  the  halogens.  It  does  not  rust, 
but  when  heated  it  burns  in  the  air,  forming  the  trioxide  Sb203  or  a 
higher  oxide  Sb204.  When  heated  with  nitric  acid,  it  yields  the  tri- 
oxide or  antimonic  acid  (H3Sb04),  according  to  the  concentration  of 
the  nitric  acid  and  the  temperature  employed.  When  heated  with 
concentrated  sulphuric  acid,  it  forms  the  sulphate  Sb2(S04)3. 

Stibine.  —  The  antimonide  of  hydrogen  SbH3  is  formed  by  the 
action  of  zinc  and  hydrochloric  acid  on  a  soluble  compound  of  anti- 
mony. By  the  action  of  dilute,  cold  hydrochloric  acid  on  an  alloy  of 
antimony  and  magnesium  (1  :  1),  a  mixture  of  hydrogen  and  stibine 
containing  as  much  as  11.5  per  cent  (by  volume)  of  the  latter  may  be 
made.  It  is  separated  by  cooling  with  liquid  air,  and  gives  a  liquid 
boiling  at  —18°  and  freezing  at  —91.5°.  It  is  more  easily  dis- 
sociated than  arsine,  and  forms  a  deposit  of  antimony  when  a  porce- 
lain vessel  is  held  in  the  flame  or  when  the  gas  passes  through  a 
heated  tube.  The  behavior  under  this  treatment  is  in  all  respects 
similar  to  that  of  arsine  (p.  709). 

The  layers  of  arsenic  or  antimony  obtained  upon  white  porcelain 
in  Marsh's  test  (p.  709)  may  be  distinguished  readily  in  several  ways. 
The  arsenic  spots  are  brownish  in  color,  lustrous,  and  volatile.  The 
antimony  spots  are  black,  smoky -looking,  and  involatile  at  the  tempera- 
ture of  the  Bunsen  flame.  The  arsenic  spots  dissolve  in  dilute  nitric 
acid,  while  those  of  antimony  do  not.  The  arsenic  spots  dissolve  in  a 
fresh  solution  of  bleaching  powder,  producing  calcium  chloride  and 


ARSENIC,    ANTIMONY,    BISMUTH  715 

arsenic  acid,  while  those  of  antimony  are  unaffected.  The  arsenic 
spots  are  scarcely  attacked  by  a  solution  of  yellow  ammonium  sulphide, 
while  those  of  antimony  dissolve  readily,  forming  an  ammonium  sul- 
phantimoniate.  Another  distinction  between  arsine  and  stibine  is  found 
in  their  action  upon  a  solution  of  nitrate  of  silver.  Stibine  precipi- 
tates a  silver  antimonide  Ag3Sb,  and  none  of  the  antimony  remains  in 
the  solution.  Arsine,  on  the  contrary,  precipitates  metallic  silver, 
while  arsenious  acid  remains  in  the  solution. 

Antimony  Halides.  —  The  halogen  compounds  of  antimony  which 
are  known  include  the  trichloride,  a  solid  melting  at  73° ;  the  penta- 
chloride,  a  liquid ;  the  tribromide,  tri-iodide,  trifluoride,  and  pentafluo- 
ride. 

Antimony  trichloride  SbCLj  is  made  by  direct  union  of  the  elements. 
It  forms  large,  soft  crystals,  and  used  to  be  named  "  butter  of  anti- 
mony." When  treated  with  water,  it  forms  a  white,  opaque,  insoluble 
basic  salt.  When  little  water  is  used,  the  product  is  antimony  oxy- 
chloride  : 

SbCl3  +  H20  <=>  SbOClJt  +  2HC1. 

With  a  large  amount  of  water,  a  greater  proportion  of  the  chlorine  is 
removed,  and  Sb405Cl2  ( =  2SbOCl,Sb203)  remains.  With  boiling  water 
the  oxide  is  finally  formed.  The  action  is  not  complete  as  long  as 
hydrochloric  acid  is  present.  It  may  therefore  be  reversed,  so  that  on 
addition  of  hydrochloric  acid  to  the  mixture,  a  clear  solution  of  the  tri- 
chloride is  re-formed.  If  the  concentration  of  the  acid  is  once  more 
reduced  by  dilution  with  water,  the  oxychloride  is  again  precipitated. 

The  pentachloride  is  f 6rmed  by  leading  chlorine  over  the  trichloride. 
It  is  a  liquid  (b.-p.  140°)  which  fumes  strongly  in  the  air,  being  hydro- 
lyzed  by  the  moisture. 

Oxides  of  Antimony.  — Three  oxides  are  known:  antimony  tri- 
oxide  Sb203,  antimony  pentoxide  Sb206,  and  an  intermediate  oxide 
Sb204.  The  trioxide  is  obtained  by  oxidizing  antimony  with  nitric  acid, 
or  by  combustion  of  antimony  with  a  limited  supply  of  oxygen.  It 
is  a  white  substance,  insoluble  in  water.  It  is  in  the  main  a  basic 
oxide,  interacting  with  many  acids  to  form  salts  of  antimony.  The 
pentoxide  is  a  yellow,  amorphous  substance,  obtained  by  heating  anti- 
monic  acid.  It  combines  with  bases  to  form  salts,  and  is  therefore  an 
acid-forming  oxide  exclusively.  The  intermediate  compound  SbaO4  is 


716  INORGANIC  CHEMISTRY 

formed  by  heating  antimony  or  the  trioxide  in  oxygen.  It  is  the  most 
stable  of  the  three  oxides.  It  is  neither  acid-  nor  base-forming,  and 
may  be  antimoniate  of  antimony  (SbSb04). 

The  hydrated  trioxide  Sb(OH)3  may  be  obtained  as  a  white  pre- 
cipitate by  adding  dilute  sulphuric  acid  to  tartar-emetic  (see  below). 
It  is  insoluble,  and  easily  loses  water,  giving  the  trioxide. 

The  trioxide  interacts  with  potassium  and  sodium  hydroxides, 
forming  soluble  antimonites,  but  the  latter  are  much  hydrolyzed  by 
water,  and  cannot  be  isolated  in  solid  form. 

Salts  of  Antimony.  — The  nitrate    Sb(N03)3   and    the    sulphate 

Sb2(S04)s  are  made  by  the  interaction  of  the  trioxide  with  nitric  and 
sulphuric  acids.  They  are  hydrolyzed  by  water,  giving  basic  salts, 
such  as  (SbO)2S04  (  =  Sb202S04),  which,  like  SbOCl,  are  derived  from 
the  hydroxide  SbO(OH).  When  the  trioxide  is  heated  with  a  solution 
of  potassium  bitartrate  KHC4H4O6,  a  basic  salt  K(SbO)C4H406,  known 
as  tartar-emetic,  is  formed.  This  is  a  white,  crystalline  substance 
which  is  soluble  in  water  and  is  used  in  medicine.  The  univalent 
group  SbO1  is  known  as  antimonyl,  and  the  above  mentioned  basic 
compounds  are  often  called  antimonyl  sulphate,  etc. 

Antimonic  Acid.  —  By  vigorous  oxidation  of  antimony  with  nitric 
acid,  or  by  decomposing  the  pentachloride  completely  with  water,  a 
white,  insoluble  substance  of  the  approximate  composition  H3Sb04  is 
obtained.  This  substance  interacts  with  caustic  potash  and  passes  into 
solution.  But  the  salts  which  have  been  made  are  pyro-  and  met- 
antimoniates.  Thus,  when  antimony  is  fused  with  niter,  potassium 
metantimoniate  KSb03  is  formed.  When  dissolved,  this  salt  takes  up 
water,  and  forms  a  solution  of  the  acid  pyroantimoniate : 

2KSb03  +  H20  -*  K2H2Sb207. 

If  this  is  added  to  a  strong  solution  of  a  salt  of  sodium,  a  sodium 
pyroantimoniate  is  thrown  down,  Na2H2Sb207.  The  same  insoluble 
body,  almost  the  only  salt  of  sodium  which  deserves  this  name,  is 
formed  also  by  direct  action  of  sodium  hydroxide  upon  antimonic 
acid. 

Sulphides  of  Antimony.  —  There  are  two  sulphides,  the  trisul- 
phide  Sb2Ss  and  the  pentasulphide  Sb2S5.  The  trisulphide  is  found  in 
nature  as  the  black,  crystalline  stibnite.  By  the  action  of  hydrogen 


ARSENIC,    ANTIMONY,    BISMUTH  717 

sulphide  upon  solutions  of  salts  of  antimony,  the  trisulphide  is  precipi- 
tated as  an  orange-red  powder,  which,  however,  after  having  been 
melted,  assumes  the  appearance  of  stibnite  : 

2SbCl3  +  3H2S  <=±  Sb2S3  J+  6HC1. 

The  antimony  trisulphide  is  decomposed,  and  the  above  action  is  re- 
versed, by  concentrated  hydrochloric  acid.  Like  cadmium  sulphide, 
this  substance  is  formed  only  when  the  acid  present  is  dilute. 

The  pentasulphide  is  obtained  by  the  decomposition  of  sulphanti- 
moniates  (see  below).  In  appearance  it  resembles  the  trisulphide  and, 
when  heated,  it  decomposes  very  readily  into  this  substance  and  free 
sulphur.  It  is  used  for  vulcanizing  rubber. 

Sulphantimonites  and  Sulphantimoniates.  —  The  behavior  of 
the  sulphides  of  antimony  towards  solutions  of  the  alkali  sulphides  is 
very  similar  to  that  of  the  sulphides  of  arsenic  (p.  712).  The  tri- 
sulphide dissolves  in  colorless  ammonium  sulphide  with  difficulty, 
forming  an  unstable  ammonium  sulphantimonite  : 

Sb2S3  +  3(NH4)2S  -*  2(NH4)8SbS3. 

With  the  pentasulphide  or  with  yellow  ammonium  sulphide  the  action 
takes  place  more  readily  and  ammonium  sulphantimoniate  is  formed : 

Sb2S5  +  3(NH4)2S  ->  2(NH4)3SbS4, 
Sb2S3  +  3(NH4)2S  +  2S-»2(NH4)3SbS4. 

The  most  familiar  substance  of  this  class  is  Schlippe's  salt  Na3SbS4, 
9H20.  Pyrargyrite  Ag3SbS3  (p.  626)  is  a  native  sulphantimonite. 

When  acids  are  added  to  solutions  of  sulphantimoniates,  the  sul- 
phantimonic  acid  which  is  liberated  decomposes,  and  antimony  penta- 
sulphide is  thrown  down  (see  under  Arsenic,  p.  713). 

BISMUTH. 

The  Chemical  Relations  of  the  Element.  —  Bismuth  forms  no 
compound  with  hydrogen.  Its  compounds  with  the  halogens  are  of 
the  form  BiX8  and  are  hydrolyzed  by  water  giving  basic  salts.  The 
oxide  Bi203  is  basic,  and,  although  an  oxide  Bi206  is  known,  it  is  not 
acidic.  Bismuth  gives  a  carbonate,  nitrate,  sulphate,  phosphate,  and 
other  salts,  in  all  of  which  it  acts  as  a  trivalent  element.  It  forms  no 
soluble  complex  sulphides. 


718  INORGANIC   CHEMISTRY 

Occurrence  and  Physical  Properties.  —  This  element  is  found 
free  in  nature,  and  also  to  some  extent  as  trioxide  Bi203  and  trisul- 
phide  BijSg.  It  is  a  shining,  brittle  metal  with  a  reddish  tinge.  It 
melts  at  about  270°.  When  converted  into  vapor,  its  density  at  1600- 
1700°  is  somewhat  less  than  that  corresponding  to  the  formula  Bi2. 

Mixtures  of  bismuth  with  other  metals  of  low  melting-point  fuse 
at  lower  temperatures  than  do  the  separate  metals.  This  is  another 
illustration  of  the  fact  that  a  solution  melts  at  a  lower  temperature  than 
the  pure  solvent  (p.  532).  Thus,  Wood's  metal,  containing  bismuth 
(m.-p.  270°)  4  parts,  lead  (m.-p.  326°)  2  parts,  tin  (m.-p.  233°)  1  part, 
and  cadmium  (m.-p.  320°)  1  part,  melts  at  60.5°,  considerably  below  the 
boiling-point  of  water.  Similar  alloys  are  used  for  safety  plugs  in 
steam-boilers,  and,  in  the  chemical  laboratory,  for  filling  baths  in  which 
a  uniform  temperature  higher  than  100°  is  to  be  maintained. 

Chemical  Properties.  —  Bismuth  does  not  tarnish,  but  when 
heated  strongly  in  the  air  it  burns  to  form  the  trioxide.  With  the 
halogens  it  forms  a  fluoride,  a  bromide,  and  an  iodide,  in  all  of  which  the 
element  is  trivalent.  When  the  metal  is  treated  with  oxygen  acids,  or 
the  trioxide  with  any  acids,  salts  are  produced. 

Oxides.  —  In  addition  to  the  basic  trioxide,  which  is  a  yellow 
powder  obtained  by  direct  oxidation  of  the  metal  or  by  ignition  of  the 
nitrate,  three  other  oxides  are  known  —  BiO,  Bi204,  and  Bi205.  None  of 
these,  however,  is  acid-forming,  or  gives  corresponding  salts  when 
treated  with  acids. 

Salts  of  Bismuth.  —  The  salts  of  bismuth,  when  dissolved  in 
water,  are  decomposed  in  a  manner  which  recalls  the  behavior  of  the 
compounds  of  antimony.  The  products  are  insoluble  basic  salts,  and 
the  actions  are  reversible,  the  basic  salts  being  redissolved  by  addition 
of  an  excess  of  the  acid.  In  the  case  of  the  chloride  BiCl3,  H20  and  the 
nitrate  Bi(NO3)s,5H2O,  the  actions  taking  place  are : 

BiCl3  +  2H20  <=±  Bi(OH)2Cl  +  2HC1, 
Bi(N03)3  +  2H30  -+  Bi(OH)2N03  +  2HN03. 

The  former  of  these  products,  when  dried,  loses  a  molecule  of  water, 
giving  the  oxychloride  BiOCl.  The  oxynitrate  of  bismuth  is  much 
used  in  medicine  under  the  name  of  "subnitrate  of  bismuth." 


ARSENIC,    ANTIMONY,    BISMUTH  719 

It  will  be  seen  that,  although  bismuth  forms  a  colorless  ion  Bi***, 
and  is  in  this  respect  a  metal  in  the  chemical  sense  of  the  term,  yet, 
like  many  other  metals,  it  is  related  to  the  non-metals  inasmuch  as  its 
salts  are  at  least  partially  hydrolyzed  by  water. 

Bismuth  forms  a  trisulphide  Bi2S8,  which  may  be  obtained  by 
direct  union  of  the  elements,  or  by  the  action  of  hydrogen  sulphide 
upon  solutions  of  bismuth  salts  : 

2BiCl8  +  3H2S  <r±  Bi2S,i  +  6HC1. 

It  is  a  brownish-black,  insoluble  substance,  but  on  addition  of  much 
acid  the  above  action  is  reversed.  This  sulphide  is  not  affected  by 
solutions  of  ammonium  sulphide  or  of  potassium  sulphide.  It  differs, 
therefore,  markedly  from  the  sulphides  of  arsenic  and  antimony  in  its 
behavior. 

THE  FAMILY  AS  A  WHOLE. 

When  we  compare  the  elements  of  this  group,  taking  nitrogen  as 
the  first  of  the  family  in  spite  of  the  fact  that  it  is  somewhat  less 
closely  related  to  the  other  members  than  they  are  to  one  another,  we 
find  an  admirable  illustration  of  the  general  principles  which  the 
periodic  system  presents. 

The  elements  themselves  change  progressively  in  physical  proper- 
ties as  the  atomic  weight  increases.  Nitrogen  is  a  gas  which  with 
sufficient  cooling  yields  a  white  solid,  phosphorus  an  almost  white,  or 
a  red  solid,  and  arsenic,  antimony,  and  bismuth  are  metallic  in 
appearance.  The  first  combines  directly  with  hydrogen,  the  next 
three  give  hydrides  indirectly,  and  the  last  does  not  unite  with 
hydrogen  at  all.  The  hydride  of  nitrogen  combines  with  water 
to  form  a  base,  while  the  other  hydrides  show  no  such  tendency. 
Ammonia  unites  with  all  acids,  including  those  of  the  halogens,  to 
form  salts ;  phosphine  with  the  hydrogen  halides  only ;  the  others  do 
not  combine  with  acids  at  all.  As  regards  their  metallic  properties, 
in  the  chemical  sense,  nitrogen  and  phosphorus  do  not  by  themselves 
form  positive  ions,  and  furnish  us  therefore  with  no  salts  whatever. 
Arsenic  gives  a  trivalent  positive  ion,  which  is  found  in  solutions  of 
the  halides  only.  It  forms  no  normal  sulphates,  nitrates,  or  other 
salts.  Antimony  and  bismuth  both  give  trivalent  positive  ions.  The 
sulphates,  nitrates,  etc.,  of  antimony,  however,  are  readily  decomposed 
by  water  with  precipitation  of  the  hydroxide.  The  salts  of  bismuth, 


720  INORGANIC   CHEMISTRY 

on  the  other  hand,  do  not  readily  give  the  pure  hydroxide  with  water, 
although  they  are  easily  hydrolyzed  to  basic  salts. 

The  halogen  compounds  of  nitrogen  and  phosphorus  are  completely 
hydrolyzed  by  water,  and  do  not  exist  when  any  water  is  present,  even 
when  excess  of  the  halogen  acid  is  used.  The  halogen  compounds  of 
arsenic  are  completely  hydrolyzed  by  cold  water,  but  exist  in  solution 
in  presence  of  excess  of  the  acids.  The  halogen  compounds  of  anti- 
mony and  bismuth  are  incompletely  hydrolyzed  by  cold  water. 

Nitrogen  forms  five  different  oxides,  a  larger  number  than  that 
furnished  by  any  of  the  other  elements  of  this  group.  The  oxides 
upon  which  we  naturally  fix  our  attention  are  the  trioxide  and  pent- 
oxide  in  each  case.  With  nitrogen  these  are  acid-forming,  being  the 
anhydrides  of  nitric  and  nitrous  acids.  With  phosphorus  the  trioxide 
and  the  pentoxide  are  anhydrides  of  acids.  With  arsenic  the  trioxide 
is  basic  towards  the  halogen  acids,  and  is  the  first  example  of  a  basic 
oxide  which  we  encounter  in  this  group.  The  pentoxide,  however,  is 
acid-forming.  The  trioxide  of  antimony  is  mainly  base-forming, 
although  it  is  feebly  acid-forming  also.  The  pentoxide  is  acid-forming. 
The  trioxide  of  bismuth  is  base-forming  exclusively,  and  the  pentoxide 
has  no  derivatives. 

If  the  chemical%  actions  which  these  elements  and  their  compounds 
undergo  were  examined  more  closely,  this  progressive  change  in  the 
properties  with  change  in  atomic  weight  could  be  developed  much 
further.  We  have  said  sufficient,  however,  to  show  that  when  the 
periodic  law  is  borne  in  mind  it  furnishes  valuable  aid  in  systematizing 
the  chemistry  of  a  group  like  this. 

Analytical  Reactions  of  Arsenic,  Antimony,  and  Bismuth.  — 

The  ions  which  are  most  frequently  encountered  are  As"',  Sb"",  Bi***, 
As04'",  and  As03'".  The  first  three,  with  hydrogen  sulphide,  give 
colored  sulphides  which  are  not  affected  by  dilute  acids.  The  sul- 
phides of  arsenic  and  antimony  are  separable  from  the  sulphide  of 
bismuth  by  solution  in  yellow  ammonium  sulphide.  The  ion  of  the 
arsenates  AsO/"  is  identified  by  its  interaction  with  salts  of  silver  and 
the -formation  of  MgNH4As04,  while  that  of  the  arsenites  As08"'  is  rec- 
ognized by  its  reducing  power.  By  means  of  Marsh's  test  and  the  facts 
describedr  under  it  (p.  714),  the  presence  of  traces  of  compounds  of 
arsenic  and  antimony  may  be  recognized  and  the  elements  may  be 
distinguished.  Oxygen  compounds  of  arsenic,  when  heated  with  car- 
bon, give  a  volatile,  metallic-looking  deposit  of  arsenic. 


ARSENIC,    ANTIMONY,    BISMUTH  721 

VANADIUM,  COLUMBIUM,  TANTALUM. 

Of  these  elements,  vanadium  is  less  uncommon  than  the  others.  It 
is  found  in  rather  complex  compounds.  When  these  are  heated  with 
soda  and  sodium  nitrate,  sodium  vanadate  is  formed,  and  can  be  ex- 
tracted with  water.  Solid  ammonium  chloride  is  added  to  the  solution, 
and  ammonium  metavanadate  NH4V03,  which  is  less  soluble  in  solu- 
tions of  salts  of  ammonium  (cf.  p.  584)  than  in  water,  appears  in  the 
form  of  yellow  crystals.  When  this  salt  is  heated,  vanadic  anhydride 
V205,  a  yellowish-red  powder,  remains.  This  oxide  interacts  with  bases 
giving  vanadates,  of  which  the  most  stable  are  the  metavanadates. 
The  element  forms  several  chlorides,  such  as  VCLj,  VC13,  VC14,  VOC18, 
and  five  oxides,  V20,  VO,  V203,  V02,  and  V205.  The  element  has  very 
feeble  base-forming  properties,  and  gives  only  a  few  unstable  salts. 

Columbium  (or  niobium)  and  tantalum  likewise  possess  feebly  base- 
forming  properties,  their  chief  compounds  being  the  columbates  and 
tantalates. 

Exercises.  —  1.  How  do  you  account  for  the  fact  that  the  molec- 
ular weight  of  arsenic  at  644°  is  not  exactly  300,  and  why  is  308.4 
-r-  4  not  accepted  as  the  atomic  weight  ?  Could  the  atomic  weight 
be  found  from  the  determination  of  the  vapor  density  of  the  free  ele- 
ment alone? 

2.  What  should  you  expect  to  be  the  interaction  of  arsine  with 
concentrated  nitric  acid  ? 

3.  Formulate  the  series  of   changes  involved   in  the  solution  of 
arsenic  trioxide   and  the   interaction  of   hydrochloric  acid  with  the 
arsenious  acid  so  formed  (cf.  p.  371). 

4.  What  is  the  full  significance  of  the  fact  that  arsenic  penta- 
sulphide  may  be  precipitated  by  hydrogen  sulphide  from  a  solution  of 
arsenic  acid  in  hydrochloric  acid  ? 

5.  To  what  classes  of  chemical  changes   do   the  interactions  of 
arsenious  sulphide  and  antimony  trisulphide  with  yellow  ammonium 
sulphide  belong? 

6.  Construct  equations  showing  the  interaction  of  (a)  concentrated 
sulphuric  acid  and  antimony,  (b)  arsenic  and  bleaching-powder  solu- 
tion, (c)  antimony  and  yellow  ammonium  sulphide,  (d)  silver  nitrate 
and  stibine,  (e)  silver  nitrate  and  arsine,  (/)  concentrated  nitric  acid 
and  antimony,  (g)  acids  and  ammonium  orthosulphantimoniate. 

7.  How  should  you  set  about  making  Schlippe's  salt  ? 


CHAPTER   XLII 
THE  CHROMIUM   FAMILY.     RADIUM 

THE  chromium  (Cr,  at.  wt.  52.1)  family  includes  molybdenum  (Mo, 
at.  wt.  96),  tungsten  (W,  at.  wt.  184),  and  uranium  (TJ,  at.  wt.  238.5), 
and  occupies  the  seventh  column  of  the  periodic  table  along  with  the 
sulphur  and  selenium  family. 

The  Chemical  Relations  of  the  Family.  —  The  features  which 
are  common  to  the  four  elements  are  also  those  which  affiliate  them 
most  closely  with  their  neighbors  on  the  right  side  of  the  column. 
They  yield  oxides  of  the  forms  Cr03,  Mo08,  W08,  and  U03,  which,  like 
S08,  are  acid  anhydrides,  and  show  the  elements  to  be  sexivalent. 
They  give  also  acids  of  the  form  H2X04,  corresponding  to  sulphuric 
acid,  whose  salts  resemble  the  sulphates.  Thus,  sodium  chromate 
Na-jCrO^  10H20  is  isomorphous  with  Glauber's  salt  (p.  576),  and  potas- 
sium chromate  K2Cr04  with  potassium  sulphate. 

Aside  from  the  chromates,  the  first  element  forms  also  two  basic 
hydroxides  Cr(OH)2  and  Cr(OH)8,  from  which  the  numerous  chromous 
(Cr")  and  chromic  (Cr***)  salts  are  derived.  Uranium  forms  a  dioxide 
UO2,  to  which  correspond  the  uranous  salts  like  U(S04)2,  but  the  most 
familiar  salts  of  this  metal  are  basic  salts  of  the  oxide  U08,  and  have 
the  form  U02(N08)2.  Molybdenum  and  tungsten  are  not  base-forming 
elements. 

CHROMIUM. 

The  Chemical  Relations  of  the  Element.  —  Chromium  gives 
four  classes  of  compounds,  and  most  of  them  are  colored  substances 
(Gk.  xpfy"1?  color).  The  chromates  are  derived  from  chromic  acid 
H2CrO4,  which,  however,  is  itself  unstable,  and  leaves  the  anhydride 
Cr08  when  its  solution  is  evaporated.  The  oxide  and  hydroxide  in 
which  the  element  is  trivalent,  namely  Cr203  and  Cr(OH)3,  are  weakly 
basic  and  still  more  weakly  acidic.  Hence  we  have  chromic  salts  such 
as  CrCl3  and  Cr2(S04)3  which  are  somewhat  hydrolyzed,  but  no  carbon- 
ate, and  no  sulphide  which  is  stable  in  water.  The  compounds  in 
which  the  same  hydroxide  acts  as  an  acid  are  the  chromites,  and  are 

722 


THE   CHROMIUM   FAMILY.     RADIUM  723 

derived  from  the  less  completely  hydrated  form  of  the  oxide  CrO(OH). 
Potassium  chromite  K.Cr02  is  more  easily  hydrolyzed,  however,  than  is 
potassium  zincate  or  potassium  aluminate.  Finally,  the  chromous  salts 
such  as  CrCl2  and  CrS04  correspond  to  chromous  hydroxide  Cr(OH)2 
in  which  the  element  is  bivalent.  This  hydroxide  is  more  distinctly 
basic  than  is  chromic  hydroxide,  and  forms  a  carbonate  and  sulphide 
which  can  be  precipitated  in  aqueous  solution.  The  chromous  salts 
resemble  the  stannous  and  ferrous  (q.v.)  salts  in  being  easily  oxidized 
by  the  air. 

Occurrence  and  Isolation.  —  Chromium  is  found  chiefly  in 
ferrous  chromite  Fe(Cr02)2,  which  constitutes  the  mineral  chromite, 
and  in  crocoisite  PbCr04,  which  is  chromate  of  lead.  The  metal  may 
be  made  by  heating  chromic  oxide  with  carbon  in  the  electric  furnace, 
or,  more  easily,  by  reduction  of  the  oxide  with  aluminium  filings  by 
Goldschmidt's  method  (p.  540). 

Physical  and  Chemical  Properties.  —  Chromium  is  steel-gray 
in  color,  very  hard,  and  extremely  infusible.  It  does  not  tarnish,  but 
when  heated  it  burns  in  oxygen,  giving  the  green  chromic  oxide  Cr203. 
It  seems  to  exist  in  two  states,  an  active  and  a  passive  one,  the  relations 
of  which  are  still  somewhat  obscure.  A  fragment  which  has  been  made 
by  the  Goldschmidt  method,  or  has  been  dipped  in  nitric  acid,  is  passive, 
and  does  not  displace  hydrogen  from  hydrochloric  acid.  When,  how- 
ever, the  specimen  is  warmed  with  this  acid,  it  begins  to  interact,  and 
thereafter  behaves  as  if  it  lay  between  zinc  and  cadmium  in  the  elec- 
tromotive series.  If  left  in  the  air,  it  slowly  becomes  inactive  again. 
Tin  and  iron  with  hydrochloric  acid  form  stannous  and  ferrous  chloride 
respectively,  because  the  higher  chlorides,  if  present,  would  be  reduced 
by  the  nascent  hydrogen.  Here,  for  the  same  reason,  chromous 
chloride  and  not  chromic  chloride  is  formed : 

Cr  +  2HC1  -4  CrCl2  +  H2  or  Cr  +  2H"  -»  Cr"+  H2. 

DERIVATIVES  OF  CHROMIC  ACID. 

Potassium  Chromate.  —  This  and  the  sodium  salt,  or  rather  the 
corresponding  dichromates  (see  below),  are  made  directly  from  chro- 
mite, and  form  the  starting-point  in  the  preparation  of  the  other  com- 
pounds of  chromium.  The  finely  powdered  mineral  is  mixed  with 
potash  and  limestone,  and  roasted.  The  lime  is  employed  chiefly  to 


724  INORGANIC   CHEMISTRY 

keep  the  mass  porous  and  accessible  to  the  oxygen  of  the  air,  the 
potassium  compounds  being  easily  fusible : 

4Fe(Cr02)2  +  8K2CO8  +  702  ->  2Fe208  +  8K2Cr04  +  8C02. 

The  iron  is  oxidized  to  ferric  oxide,  and  the  chromium  passes  from  the 
state  of  chromic  oxide  in  the  chromite  (FeO,Cr208)  to  that  of  chromic 
anhydride  in  the  potassium  chromate  (K20,Cr03).  Thus,  more  insight 
is  given  into  the  nature  of  the  action  by  the  equation 

4(FeO,Cr203)  +  8(K,0,C02)  +  702  ~>  2Fe203  +  8(K20,O03)  +  '8C02. 

The  cinder  is  treated  with  hot  potassium  sulphate  solution.  This 
interacts  with  the  calcium  chromate,  which  is  formed  at  the  same 
time,  giving  insoluble  calcium  sulphate : 

CaCr04  +  K2S04  <=>  CaS04  J+  K2CrO4. 

The  whole  of  the  potassium  chromate  goes  into  solution. 

Potassium  chromate  is  pale-yellow  in  color,  rhombic  in  form  (iso- 
morphous  with  potassium  sulphate),  and  is  very  soluble  in  water 
(61 :  100  at  10°). 

Sodium  chromate  is  made  by  using  sodium  carbonate  in  the 
process  just  described. 

The  Dichromates.  —  When  a  solution  of  potassium  sulphate  is 
mixed  with  an  equivalent  amount  of  sulphuric  acid,  potassium  bisul- 
phate  is  obtainable  by  evaporation  :  K2S04  -f-  H2S04  — >  2KHS04.  The 
dry  acid  salt,  when  heated,  loses  water  (p.  388),  giving  the  pyrosulphate 
(or  disulphate) :  2KHS04  +±  K2S207  +  H20,  but  the  latter,  when  redis- 
solved,  returns  to  the  condition  of  acid  sulphate.  Now,  when  an  acid 
is  added  to  a  chromate  we  should  expect  the  chromic  acid  H2Cr04,  thus 
liberated,  to  interact,  giving  an  acid  chromate  (say,  KHCr04).  No  acid 
chromates  are  known,  however,  and  instead  of  them,  pyrochromates  or 
dichromates  are  produced,  with  elimination  of  water.  In  other  words, 
the  second  of  the  above  actions  is  not  appreciably  reversible  when 
chromates  are  in  question  : 


K2Cr04  H 
K2Cr04  ( 

h  H2S04 
+  H2Cr04) 

-»(H2Cr04). 
-»  K2Cr207  4 

-h  K2S04 
•H20 

2K2Cr04  +  H2S04     ->  K2Cr207  +  H2O  4-  K2SO4  (1) 


THE   CHROMIUM  FAMILY.     RADIUM  725 

In  terms  of  the  ionic  hypothesis,  S2O7"  is  unstable  in  water,  and  in. 
teracts  with  it,  giving  hydrion  and  sulphanion,  while  Cr2O7"  is  stable 
in  water  and  is  formed  from  the  interaction  of  hydrion  and  chromanion : 

S207"    +  H20  fc?  2H'  +  2S04", 

Cr207"  +  H20  <=;  2H-  -f  2CrO4".  (2) 

The  dichromates  of  potassium  and  sodium  are  made  by  adding 
sulphuric  acid  to  the  crude  solution  of  the  chromate  obtained  from 
chromite  (p.  724).  They  crystallize  when  the  liquid  cools,  and  the 
mother-liquor,  containing  the  potassium  sulphate  and  undeposited 
dichromate,  is  used  for  extracting  a  fresh  portion  of  cinder.  As  the 
dichromates  are  much  less  soluble  than  the  chromates,  they  crystallize 
from  less  concentrated  solutions,  and  can  therefore  be  obtained  in 
purer  condition.  For  this  reason  the  extract  is  always  treated  for 
dichromate. 

Potassium  dichromate  K^Cr^  (or  K^CrO^CrOg)  crystallizes  in 
asymmetric  tables  of  orange-red  color.  Its  solubility  in  water  is  8 : 
100  at  10°  and  12.5  : 100  at  20°.  Sodium  dichromate  Na2Cr207, 2H20 
forms  red  crystals  also,  and  its  solubility  is  109 : 100  at  15°.  This  salt 
is  now  cheaper  than  potassium  dichromate,  and  has  largely  displaced 
the  latter  for  commercial  purposes. 

By  treatment  of  the  chromates  with  larger  amounts  of  free  acid, 
other  polychromates  are  formed.  Thus,  with  increasing  amounts  of 
nitric  acid,  ammonium  chromate  gives  first  the  dichromate  (NH4)2Cr207, 
which  may  be  written  (NH4)2Cr04,Cr03,  then  the  trichromate  (NH4)2 
Cr04, 2Cr03,  and  even  the  tetrachromate  (NH4)2CrO4, 3Cr08,  all  of  which 
are  red  crystalline  substances. 

Chemical  Properties  of  the  Dichromates.  —  1.  When  concen- 
trated sulphuric  acid  is  added  to  a  strong  solution  of  a  dichromate  (or 
chromate),  chromic  anhydride  separates  in  red  needles  : 

Na2Cr207  +  H2S04  ->  Na2S04  +  H20  +  2Cr08. 

2.  Although  a  dichromate  lacks  the  hydrogen,  it  is  essentially  of 
the  nature  of  an  acid  salt,  just  as  SbOCl  lacks  hydroxyl,  but  is  essen- 
tially a  basic  salt.  Hence,  when  potassium  hydroxide  is  added  to  a 
solution  of  potassium  dichromate,  potassium  chromate  is  formed : 

K2020T  +  2KOH  ->  2K2Cr04  +  H20. 


726  INORGANIC   CHEMISTRY 

The  solution  changes  from  red  to  yellow,  and  the  chromate  is  obtained 
by  evaporation.  It  is  in  this  way  that  pure  alkali  chromates  are 
made. 

3.  By  addition  of  potassium  dichromate  to  a  solution  of  a  salt  of  a 
metal  whose  chromate  is  insoluble,  the  chromate  and  not  the  dichro- 
mate is  precipitated.  This  is  in  consequence  of  the  fact  that  there  is 
always  a  little  hydrion  and  Cr04"  (equation  (2)  above)  in  the  solution 
of  the  dichromate  : 


2Ba(N03)2  4-  K^O,  +  H2O  «=*  2BaCr04|  +  2KN03  +  2HN03. 

Being  essentially  an  acid  salt,  the  dichromate  produces  a  salt  and  an 
acid,  as  any  acid  salt  would  do.     For  example  : 


Ba(NO8)2  +  KHS04  <±  BaSOJ  +  KN08  +  HN03. 

Soluble  chromates  are,  naturally,  equally  good  precipitants  of  insoluble 
ones. 

4.  The  dichromates  of  potassium  and  sodium  melt  when  heated,  and, 
at  a  white  heat,  decompose,  giving  the  chromate,  chromic  oxide,  and  free 
oxygen.  To  make  the  equation,  we  note  that  the  dichromate,  for 
example  KjjCrjjO^  consists  of  K2Cr04  -f-  Cr03,  and  the  latter,  if  alone, 
will  decompose  thus  :  2Cr03  —  >  O2O3  -f  30.  Since  the  product  must 
contain  a  multiple  of  O2,  the  equation  is  : 


4K2Cr04  +  2O2O3  +  30 


2. 


5.  With  free  acids  the  dichromates  give  powerful  oxidizing  mix- 
tures, in  consequence  of  their  tendency  to  form  chromic  salts.  Since 
the  latter  correspond  to  the  oxide  O203  and  the  former  to  Cr03,  the 
passage  from  the  former  to  the  latter  must  furnish  30  for  every 
2Cr03  transformed.  In  dilute  solutions,  unless  a  body  capable  of 
being  oxidized  is  present,  no  actual  decomposition,  beyond  the  libera- 
tion of  chromic  acid  *  occurs.  When  concentrated  hydrochloric  acid  is 
used,  this  acid  itself  suffers  oxidation  : 

K20207  +   8HC1  -r*  2KC1  +  2CrCl3  +  4H2O  (  +  30) 
(30)       +  6HC1  -»  3H20  +  301,  _ 
K2Cr2O7  +  14HC1  ->  2KC1  +  2CrCl8  +  7H2O  +  SGI, 

When  sulphuric  acid  is  employed,  an  oxidizable   substance  such  as 
*  Not  shown  as  a  distinct  stage  in  the  subsequent  equations. 


THE    CHROMIUM   FAMILY.     RADIUM  727 

hydrogen  sulphide  (cf.  p.  374),  sulphurous  acid,  or  alcohol  must  be 
present : 

K2Cr207  +      4H2S04  -»  K,S04  +  Cr2(S04)8  +  4H20  ( +  30)     (1) 

(30)  +      3H2S03  -»  3H2S04  (2) 

or  (30)  +  3C2H5OH  -+  3C2H40  f  +  3H20  (2') 

[alcohol]  [aldehyde] 

In  each  case  the  usual  summation  of  (1)  and  (2),  with  omission  of  the 
30  gives  the  equation  for  the  whole  action.  When  (1)  is  dissected, 
K20,  2Cr03  giving  3S03,Cr203+30  is  found  to  be  its  essential  content. 
In  practice,  this  sort  of  action  is  used  for  the  purpose  of  making 
chromic  salts,  and  for  its  oxidizing  effects,  as  in  the  preparation  of 
aldehyde  and  in  the  dichromate  battery  (q.v.). 

6.  When  a  body  which  is  not  merely  oxidizable,  but  is  an  active 
reducing  agent,  is  employed,  the  dichromate  may  be  reduced  without 
the  addition  of  any  acid.  For  example,  when  warmed  with  ammonium 
sulphide,  a  dichromate  gives  chromic  hydroxide  and  free  sulphur : 

K2Cr207  +  3(NH4)2S  +  7H20  -»  2Cr(OH)s  +  3S  +  2KOH  +  6NH4OH. 

If  the  reducing  body  is  less  active,  the  change  may  nevertheless  take 
place  under  the  influence  of  light.  Thus,  when  paper  is  coated  with 
gelatine  containing  a  soluble  chromate  or  dichromate,  and,  after  being 
dried,  is  exposed  to  light,  chromic  oxide  is  formed  by  reduction,  and 
combines  with  the  gelatine  to  give  an  organic  compound.  This  product 
will  not  swell  up  or  dissolve  in  tepid  water  as  does  pure  gelatine. 
This  action  is  used  in  many  ways  for  purposes  of  artistic  reproduction. 
Thus,  if  the  gelatine  mixture  is  made  up  with  lampblack,  and,  after  the 
coating  has  dried,  is  covered  with  a  negative  and  exposed  to  light,  the 
parts  which  were  protected  from  illumination  may  afterwards  be 
washed  away,  while  the  "  carbon  print "  remains.  The  gelatine  layer 
can  be  transferred  to  wood  or  copper  before  washing.  When  materials 
of  different  colors  are  substituted  for  the  lampblack,  prints  of  any 
desired  tint  may  be  made  by  the  same  process. 

Insoluble  Chromates.  —  A  number  of  chromates,  formed  by  pre- 
cipitation with  a  solution  of  a  soluble  chromate  or  dichromate,  are 
familiar.  Thus,  lead  chromate  PbCr04  is  used  as  a  yellow  pigment. 
By  treatment  with  lime-water  it  gives  a  basic  salt  of  brilliant  red  color 
—  "  chrome-red  "  Pb2OCr04.  Salts  of  calcium  give  a  yellow,  hydrated 
calcium  chromate  CaCr04,  2H02  analogous  to  gypsum,  and,  like  it,  per- 


728  INORGANIC   CHEMISTRY 

ceptibly  soluble  in  water  (0.4  :  100  at  14°).  Barium  chromate  BaCr04 
is  also  yellow.  Being  a  salt  of  a  feeble  acid,  it  interacts  with  active 
acids,  and  passes  into  solution.  Like  calcium  oxalate  (cf.  p.  599),  it  is 
not  soluble  enough  to  be  attacked  by  acetic  acid.  Strontium  chromate, 
however,  is  soluble  in  acetic  acid.  Silver  chromate  is  red,  and  inter- 
acts easily  with  acids.  It  will  be  observed  that  there  is  a  close  corre- 
spondence between  the  relative  solubilities  of  the  chromates  and  the 
sulphates. 

Chromyl  Chloride.  — This  compound  corresponds  to  sulphuryl 
chloride  SO-jCJ^,  and  is  made  by  distilling  a  dichromate  with  a  chloride 
and  concentrated  sulphuric  acid  : 

K2Cr207  +  4KC1  +  3H2S04  -+  2Cr02Cl2  +  3K2S04  +  3H2O. 

The  hydrochloric  acid  liberated  from  the  chloride  may  be  supposed  to 
interact  with  chromic  acid  from  the  dichromate  : 

Cr02(OH)2  +  2HC1  -+  CrC^CL,  +  2H20. 

Chromyl  chloride  is  a  red  liquid,  boiling  at  118°.  It  fumes  strongly  in 
moist  air,  being  hydrolyzed  by  water.  This  action  is  the  reverse  of 
that  shown  in  the  last  equation.  No  corresponding  bromine  and  iodine 
compounds  are  known ;  and  when  a  bromide  or  iodide  is  treated  as  de- 
scribed above,  the  halogens  are  liberated  by  oxidation,  and  no  volatile 
compound  of  chromium  appears.  Hence,  when  an  unknown  halide  is 
mixed  with  potassium  dichromate  and  sulphuric  acid  and  distilled,  and 
the  vapors  are  caught  in  ammonium  hydroxide,  the  finding  of  a  chromate 
in  the  distillate  demonstrates  the  existence  of  a  chloride  in  the  original 
substance : 

Cr02Cl2  +  4NH4OH  -+  (NH4)2Cr04  +  2NH4C1  +  2H20. 

This  action  is  used  as  a  test  for  the  presence  of  traces  of  chlorides  in 
large  amounts  of  bromides  or  iodides. 

Chromic  Anhydride.  —  The  oxide  Cr03  is  made  as  described  above 
(par.  1,  p.  725),  and  is  often  called  chromic  acid.  It  is  soluble  in 
water,  and  combines  with  the  latter  to  some  extent,  giving  dichromic 
acid  H2.Cr2O7.  In  acidified  solution  it  is  much  used  as  an  oxidizing 
agent  for  organic  substances.  It  interacts  with  acids  in  the  same  way 
as  do  the  dichromates,  giving  chromic  salts  and  furnishing  oxygen  to 


THE   CHROMIUM   FAMILY.     RADIUM  729 

the  oxidizable  body.  When  heated  by  itself,  it  loses  oxygen  readily, 
and  yields  the  green  chromic  oxide  :  4Cr08  — >  2Cr2O3  +  302. 

Perchromic  Acid.  —  When  hydrogen  peroxide  (cf.  p.  306)  is  pres- 
ent in  solution  with  free  chromic  acid,  a  deep-blue,  unstable  compound 
is  formed.  It  can  be  extracted  from  the  solution  by  means  of  ether, 
and  its  formation  is  used  as  a  delicate  test  for  hydrogen  peroxide 
(q.v.).  It  is  supposed  to  be  a  mixed  anhydride  of  hydrogen  peroxide 
with  dichromic  acid,  perhaps  H2Cr208  (cf.  p.  397). 

CHROMIC    COMPOUNDS. 

Chromic  Chloride.  —  A  hydrated  chloride  CrCl3,6H20  is  ob- 
tained by  treating  the  hydroxide  Cr(OH)8  with  hydrochloric  acid  and 
evaporating.  When  heated,  this  is  hydrolyzed,  and  chromic  oxide  re- 
mains. The  anhydrous  chloride  is  formed  by  sublimation,  as  a  mass  of 
brilliant,  reddish- violet  scales,  when  chlorine  is  led  over  a  heated  mix- 
ture of  chromic. oxide  and  carbon  (cf.  Silicon  tetrachloride) :. 

Cr203  +  3C  +  3C12  -»  2CrCl,  +  3CO. 

In  this  form  the  substance  dissolves  with  extreme  slowness,  even  in 
boiling  water,  but  in  presence  of  a  trace  of  chromous  chloride  or 
stannous  chloride  it  is  easily  soluble.  The  solution  is  green,  as  are 
all  solutions  of  chromic  salts  after  they  have  been  boiled,  but  on  stand- 
ing in  the  cold,  blue  crystals  of  CrCl3,  6H20  are  deposited.  These  give 
a  blue  solution  containing  Cr"""  -j-  301',  but  boiling  reproduces  the 
green  color.  The  green  material  is  a  basic  salt  of  an  exceptional 
nature.  It  has  lost  one  unit  of  chlorine  by  hydrolysis,  and  one  of  the 
two  others  is  not  precipitated  by  nitrate  of  silver.  The  substance  is 
supposed,  therefore,  to  contain  a  complex  cation  and  to  have  the  formula 
CrClOH.Cl. 

Chromic  Hydroxide.  —  When  ammonium  hydroxide  is  added  to 
a  solution  of  a  chromic  salt,  a  hydrated  hydroxide  of  pale-blue  color, 
Cr(OH)3,  2H20,  is  thrown  down.  This  loses  water  by  stages,  giving 
intermediate  hydroxides  such  as  Cr(OH)3  and  CrOOH,  and  finally 
Cr208.  It  interacts  with  acids,  giving  chromic  salts.  It  also  dis- 
solves in  potassium  and  sodium  hydroxides  to  form  green  solutions  of 
chromites  of  the  form  KCrO2.  When  the  solutions  of  the  alkali 
chromites  are  boiled,  the  free  hydroxide,  present  in  consequence  of 


730  INORGANIC   CHEMISTRY 

hydrolysis,  is  converted  into  a  greenish,  less  completely  hydrated,  and 
less  soluble  variety.  This  begins  to  come  out  as  a  precipitate,  and 
soon  the  whole  action  is  reversed.  Insoluble  chromites,  such  as  that 
of  iron  Fe(Cr02)2,  are  found  in  nature.  Many  of  them,  like  Zn(Cr02)2 
and  Mg(Cr02)2,  may  be  formed  by  fusing  the  oxide  of  the  metal  with 
chromic  oxide ;  the  action  being  similar  to  that  used  in  making  zinc- 
ates  (p.  648)  and  aluminates  (p.  691). 

Chromic  Oxide.  —  This  oxide  is  obtained  as  a  green,  infusible 
powder  by  heating  the  hydroxide ;  or,  more  readily,  by  heating  dry 
ammonium  dichromate ;  or  by  igniting  potassium  dichromate  with  sul- 
phur and  washing  the  potassium  sulphate  out  of  the  residue  : 

(NH4)2Cr207  -»  N2  +  4H20  +  Cr208, 
K2O207  +  S  -»  K2S04  +  Cr203. 

Chromic  oxide  is  not  affected  by  acids,  but  may  be  converted  into  the 
sulphate  by  fusion  with  potassium  bisulphate.  It  is  used  for  making 
green  paint,  and  for  giving  a  green  tint  to  glass.  When  the  oxide,  or 
any  of  the  chromic  salts,  is  fused  with  a  basic  substance  such  as  an 
alkali  carbonate,  it  passes  into  the  form  of  a  chromate,  absorbing  the 
necessary  oxygen  from  the  air.  If  an  alkali  nitrate  or  chlorate  is  added, 
the  oxidation  goes  on  more  quickly.  The  alkaline  solution  of  the 
chromites  may  be  oxidized,  for  example,  by  addition  of  chlorine  or 
bromine,  and  chromates  are  formed. 

Chromic  Sulphate,  —  This  salt  crystallizes  in  reddish-violet  crys- 
tals of  a  hydrate  Cr2(S04)3,15H20,  and  may  be  made  by  treating  the 
hydroxide  with  sulphuric  acid.  It  gives  reddish-violet,  octahedral 
crystals  of  chrome-alum  (cf.  p.  687),  K2S04,Cr2(S04)3, 24H2O,  when 
mixed  with  potassium  sulphate.  This  double  salt  is  most  easily  ob- 
tained by  reducing  potassium  dichromate  in  dilute  sulphuric  acid  by 
means  of  sulphurous  acid  (p.  727),  and  allowing  the  solution  to  crystal- 
lize. The  solution  of  the  crystals,  either  of  the  pure  sulphate  or  of 
the  alum,  is  bluish-violet  (Cr***),  but  when  boiled  becomes  green.  The 
green  compound  is  formed  by  hydrolysis  and  is  gummy  and  uncrystal- 
lizable.  It  even  yields  products  which  do  not  show  the  presence  either 
of  the  Cr"*  or  the  S04"  ion.  It  seems  to  be  formed  thus  : 

2Cr2(SO4)8  +  H20  <=»  Cr40(S04)4.SO4  +  H2S04. 


THE   CHROMIUM  FAMILY.     RADIUM  731 

The  green  materials  revert  slowly  to  the  violet  ones  by  reversal  of  the 
above  action  when  the  solution  remains  in  the  cold,  and  so  crystals  of 
the  sulphate  or  of  the  alum  are  obtainable  from  the  green  solutions. 

Chromic  Acetate.  — This  salt,  Cr(C2H802)8,  is  made  by  treating 
the  hydroxide  with  acetic  acid,  and  a  green  solution  of  it  is  used  as  a 
mordant  by  calico-printers  (of.  p.  689). 

CHROMOUS  COMPOUNDS. 

By  the  interaction  of  chromium  with  hydrochloric  acid,  or  by  re- 
ducing chromic  chloride  in  a  stream  of  hydrogen,  chromous  chloride 
CrCl2  is  formed.  The  anhydrous  salt  is  colorless,  and  its  solution  is 
blue  (Or**).  Like  stannous  chloride,  it  is  very  easily  oxidized  by  the 
air,  a  solution  of  it  containing  excess  of  hydrochloric  acid  being  used 
in  the  laboratory  to  absorb  oxygen : 

4CrCl2  +  4HC1  +  O2  ->  4CrCl3  +  2H20. 

Chromous  hydroxide  is  obtained  as  a  yellow  precipitate  when  alka- 
lies are  added  to  the  chloride.  With  sulphuric  acid  it  gives  chromous 
sulphate  CrS04,  7H20,  which  is  isomorphous  with  the  vitriols  (p.  649). 

Chromous  salts  give  with  ammonium  sulphide  a  black  precipitate 
of  chromous  sulphide,  and  with  sodium  acetate  a  red  precipitate  of 
chromous  acetate.  The  latter  is  not  very  soluble,  and  is  less  quickly 
oxidized  by  the  air  than  any  of  the  other  chromous  compounds. 

Analytical    Reactions    of    Chromium    Compounds.  —  The 

chromic  salts  give  the  bluish- violet  trichromion  Cr*",  or  the  green  com- 
plex cations,  and  may  be  recognized  in  solution  by  their  color.  The 
chromates  and  dichromates  give  the  ions  Cr04"  and  Cr207",  which  are 
yellow  and  red  respectively.  From  chromic  salts,  alkalies  and  ammo- 
nium sulphide  precipitate  the  bluish-green  hydroxide,  and  carbonates 
give  a  basic  carbonate  which  is  almost  completely  hydrolyzed  to 
hydroxide.  By  fusion  with  sodium  carbonate  and  sodium  nitrate,  they 
yield  a  yellow  bead  containing  the  chromate.  The  chromates  and 
dichromates  are  recognized  by  the  insoluble  chromates  which  they 
precipitate,  and  by  their  oxidizing  power  when  mixed  with  acids.  All 
compounds  of  chromium  give  a  green  borax  bead  containing  chromic 
borate,  and  this  bead  differs  from  that  given  by  compounds  of  copper 
(cf.  p.  626),  which  is  also  green,  in  being  unreducible. 


732  INORGANIC   CHEMISTRY 

MOLYBDENUM,  TUNGSTEN,  URANIUM. 

As  was  stated  at  the  opening  of  the  chapter,  these  elements  give 
acid  anhydrides  of  the  form  XO3,  and  acids  and  salts  of  the  form 
H2X04.  They  also  give  salts  of  the  form  H^O,  corresponding,  to 
the  dichromates.  Uranium  has  base-forming  properties  as  well. 

Molybdenum.  —  This  element  is  found  chiefly  in  wulf enite 
PbMo04  and  molybdenite  MoS2.  The  latter  resembles  black  lead 
(graphite),  and  its  appearance  suggested  the  name  of  the  element 
(Gk.  /AoAv/3&xiva,  lead).  The  molybdenite  is  converted  by  roasting  into 
molybdic  anhydride  Mo03.  When  this  is  treated  with  ammonium 
hydroxide,  or  with  sodium  hydroxide,  ammonium  molybdate  (NH4)2 
Mo04,  or  sodium  molybdate  Na2Mo04, 10H20  is  obtained.  The  metal 
itself  is  liberated  by  reducing  the  oxide  or  chloride  with  hydrogen. 
When  pure  it  resembles  wrought  iron,  and,  like  iron  (q.v.),  takes  up 
carbon  and  shows  the  phenomena  of  tempering.  The  oxides  [MoO  ?], 
Mo203,  Mo02,  and  Mo03  are  known,  but  the  lower  oxides  are  not  basic. 
The  chlorides  MosCl6,  MoCl3,  MoCl4,  and  MoCl6  have  been  made.  The 
chief  use  of  molybdenum  compounds  in  the  laboratory  is  in  testing  for 
and  estimating  phosphoric  acid.  When  a  little  of  a  phosphate  is  added 
to  a  solution  of  ammonium  molybdate  in  nitric  acid,  and  the  mixture 
is  warmed,  a  copious  yellow  precipitate  of  a  phosphomolybdate  of 
ammonium  (NH4)8P04,llMo03, 6H20  is  formed.  The  compound  is 
soluble  in  excess  of  phosphoric  acid  and  in  alkalies,  but  not  in  dilute 
mineral  acids. 

Tungsten.  —  The  minerals  scheelite  CaW04  and  wolfram  FeW04 
are  tungstates  of  calcium  and  iron  respectively.  By  fusion  of  wolfram 
with  sodium  carbonate  and  extraction  with  water,  sodium  tungstate 
Na2W04,  2H20  is  secured.  It  is  used  as  a  mordant  and  for  rendering 
muslin  fire-proof.  Acids  precipitate  tungstic  acid  H2W04,H20  from 
solutions  of  this  salt.  The  element  gives  the  oxides  W02  and  W03, 
the  latter  being  formed  by  ignition  of  tungstic  acid.  The  chlorides 
WC12,  WC14,  WC16,  and  WC16  are  known,  the  last  being  formed 
directly,  and  the  others  by  reduction.  A  hard  variety  of  steel  contains 
5  per  cent  of  tungsten. 

Uranium.  —  This  element  is  found  chiefly  in  pitchblende,  which 
contains  the  oxide  U808  along  with  smaller  amounts  of  many  other  ele- 


THE    CHROMIUM  FAMILY.     RADIUM  733 

merits.  By  roasting  the  ore  with  carbonate  and  nitrate  of  sodium,  and 
extracting  with  water,  an  impure  solution  of  sodium  uranate  NaUO4  is 
obtained.  Acids  precipitate  the  insoluble,  yellow  diuranate  Na^U^, 
6H20.  This  salt  is  used  in  making  uranium  glass,  which  shows  a 
yellowish-green  fluorescence.  The  property  is  due  to  the  fact  that 
the  wave-length  of  part  of  the  invisible,  ultra-violet  rays  of  the  sun- 
light are  lengthened,  and  a  greenish  light  is  therefore  in  excess.  The 
oxides  are  U02  a  basic  oxide,  U208,  U808  the  most  stable  oxide, 
U03  uranic  anhydride,  and  UO4  a  peroxide. 

When  the  oxide  U02  is  treated  with  acids,  it  gives  uranous  salts 
such  as  uranous  sulphate  U(S04)2/4H20.  Uranic  anhydride  and 
uranic  acid  interact  with  acids,  giving  basic  salts,  such  as  U02S04, 
3£H20,  and  U02(N03)2, 6H20,  which  are  named  uranyl  sulphate,  uranyl 
nitrate,  and  so  forth.  They  are  yellow  in  color,  with  green  fluores- 
cence. Ammonium  sulphide  throws  down  the  brown,  unstable  uranyl 
sulphide  U02S  from  their  solutions. 

RADIUM. 

The  Discovery  of  the  Element.*  —  It  was  Becquerel  who  first 
noticed  (1896)  that  all  compounds  of  uranium  gave  out  a  radiation 
capable  of  affecting  a  photographic  plate  covered  with  black,  light- 
proof  paper.  The  Becquerel  rays,  however,  required  several  days  to 
produce  a  distinct  effect.  These  rays  had  a  second,  equally  remarkable 
property.  Ordinary  air  is  an  extremely  poor  conductor  of  electricity, 
and  for  this  reason  a  well-insulated,  electrically  charged  body,  such  as 
an  electroscope,  will  retain  its  charge  for  a  long  time.  Yet  a  few 
tenths  of  a  gram  of  any  uranium  compound,  brought  within  3  or  4  cm. 
of  the  charged  body,  rendered  the  air  a  conductor,  and  the  charge  was 
quickly  lost.  The  air,  under  these  conditions,  is  said  to  be  "  ionized," 
but  the  positive  and  negative  ions  it  contains  are  probably  large 
molecular  complexes.  This  property  makes  possible  the  quantitative 
measurement  of  radio-activity,  or  the  rate  of  production  of  Becquerel 
rays.  We  simply  have  to  compare  the  times  required  for  the  discharge 
of  an  electroscope  by  different  specimens  of  radio-active  matter. 

The  radio-activity  of  every  pure  uranium  compound  is  proportional 
to  its  uranium  content.  The  ores  are,  however,  relatively  four  times 
as  active.  This  fact  led  M.  and  Mine.  Curie  to  the  discovery  that  the 

*  I  am  indebted  to  my  colleague  Professor  H.  N.  McCoy  for  the  material  of 
which  the  following  is  a  slightly  condensed  version. 


734  INORGANIC   CHEMISTRY 

pitchblende  residues,  from  which  practically  all  of  the  uranium  had 
been  extracted,  were  nevertheless  quite  active.  About  a  ton  of  the 
very  complex  residues  having  been  separated  laboriously  into  the  con- 
stituents, it  was  found  that  a  large  part  of  the  radio-activity  remained 
with  the  compound  of  barium.  The  barium  chloride,  after  being  puri- 
fied until  no  other  elements  could  be  detected  in  it  by  ordinary  chemi- 
cal tests,  was  sixty  times  as  active  as  uranium.  The  percentage  of 
chlorine  contained  in  it  was  almost  identical  with  that  in  the  ordinary 
salt,  and  it  differed  from  this  only  in  its  photo-active  and  "  ionizing  " 
powers.  By  repeated  systematic  recrystallization,  however,  a  portion 
was  separated  which  gave  distinctly  the  spectrum  of  a  new  element  and 
a  diminished  percentage  of  chlorine.  Finally,  a  product  free  from 
barium,  and  three  million  times  as  active  as  uranium,  was  secured.  The 
nature  of  the  spectrum  and  the  chemical  relations  of  the  element,  now 
named  radium,  placed  it  with  the  metals  of  the  alkaline  earths.  The 
ratio  by  weight  of  chlorine  to  radium  in  the  compound  was  35.46: 
113.25,  so  that,  on  the  assumption  that  the  element  is  bivalent,  its 
atomic  weight  is  226.5.  With  this  value  it  occupies  a  place  formerly 
vacant  in  the  periodic  table. 

Properties  of  Radium  Compounds.  —  Radium  has  not  been 
isolated,  and  the  chloride  and  bromide  are  usually  employed.  The 
enormous  photo-activity  and  ionizing  power  of  the  compound  has 
been  mentioned  above.  The  rays,  like  X-rays,  may  cause  severe 
"burns."  Many  substances,  like  zinc-blende  and  the  diamond,  phos- 
phoresce brilliantly  when  exposed  to  the  rays.  The  most  remarkable 
property  of  the  salts,  however,  is  their  constant  evolution  of  heat  in 
relatively  enormous  quantities.  One  gram  of  the  element,  in  com- 
bination, evolves  over  100  cal.  per  hour,  and  it  is  estimated  that  the 
total  amount  of  heat  spontaneously  produced  by  1  g.  would  be  about 
1010  cal.  To  produce  the  same  amount  of  heat  by  combustion,  no  less 
than  300  kg.  of  hydrogen  would  have  to  be  burned. 

The  Radiations  of  Radium  Salts.  —  The  Becquerel  rays  are 
made  up  of  three  different  radiations :  The»  a-rays,  which  produce  the 
ionization,  and  are  almost  completely  absorbed  by  a  piece  of  paper ; 
the  /8-rays,  which  produce  the  photographic  effects  and  readily  penetrate 
paper  and  even  thin  sheets  of  metal;  the  y-rays,  which  resemble 
X-rays. 

When  ammonium  carbonate  is  added  in  excess  to  a  solution  of  a 


THE   CHROMIUM   FAMILY.     RADIUM  735 

soluble  salt  of  uranium,  a  precipitate  is  formed  and  redissolves.  A 
trifling  undissolved  residue,  called  uranium-X,  however,  possesses 
itself  of  all  the  /?-ray  activity  of  the  original  material.  While  the 
activity  of  any  ordinary  uranium  compound  is  constant,  that  of  U-X 
decreases  rather  rapidly,  reaching  half  value  in  22  days,  and  disappear- 
ing entirely  in  a  few  months. 

The  uranium  freed  from  U-X  has  at  first  only  o-ray  activity,  but 
after  22  days  it  has  recovered  half  its  former  /3-ray  activity,  and 
after  a  few  months  the  whole.  The  removal  of  equal  quantities  of 
U-X  may  be  repeated,  in  this  way,  at  intervals  of  a  few  months,  any 
number  of  times.  There  is  thus  a  continuous  production  of  U-X  from 
uranium. 

It  has  been  found  that  the  /3-rays  are  identical  with  cathode  rays, 
and,  in  terms  of  the  molecular  hypothesis,  consist  of  minute  particles, 
electrons  or  corpuscles,  shot  out  with  a  velocity  approaching  that  of 
light.  The  mass  of  each  corpuscle  is  about  y^^  of  that  of  an  atom 
of  hydrogen  (cf.  p.  222),  and  bears  a  negative  charge  of  electricity 
equal  to  that  on  a  chlorine  ion  in  solution. 

The  a-rays  consist  of  particles  about  twice  as  heavy  as  a  hydrogen 
atom,  and  move  with  enormous  velocity.  Each  of  these  particles  bears 
a  positive  charge  equal  to  that  on  one  atom  of  hydrion. 

Rutherford  and  Soddy  account  for  the  facts  on  which  the  above 
hypothetical  statements  are  based  by  the  disintegration  hypothesis. 
All  atoms  are  considered  as  groups  of  minute  particles  in  a  state  of 
rapid  orbital  motion.  The  atoms  of  radio-active  elements  are  rot 
perfectly  stable,  and,  occasionally,  disintegration  occurs,  one  or  more  of 
the  particles  being  shot  out.  These  particles  constitute  the  a-  and 
/3-rays.  The  heating  effect  of  radium  salts  is  thus  due  to  the  change 
of  the  kinetic  energy  of  the  moving  particles  into  heat  when  they 
encounter  some  other  body. 

The  y-rays  probably  originate  by  the  impact  of  /3-rays,  just  as 
X-rays  are  produced  by  cathode  rays. 

The  Decay  of  an  Element.  —  Uranium  (at.  wt.  238.5)  has  the 
heaviest  known  atom.  The  disintegration  of  its  atom  gives  a-rays  and 
leaves  the  lighter  atoms  of  U-X.  In  a  similar  manner,  radium  gives 
rise  to  a  new  radio-active,  gaseous  body,  the  radium  emanation.  This, 
like  U-X,  is  not  permanent,  and  loses  half  its  activity  in  four  days.  In 
doing  so  it  produces  a  series  of  new  radio-active  substances.  These, 
with  their  times  of  decay  to  half  value,  are  as  follows  (Rutherford) : 


736  INORGANIC   CHEMISTRY 

Ra^Em— »A^B-»C-->D->E— >     F 
2000  yrs.  4  days  3min.  26min.  19min.  40  yrs.  6  days    45  days 

Radium  thus  loses  its  activity.  To  account  for  its  presence  in  the  ore, 
we  must,  therefore,  suppose  that  it  is  being  continuously  produced  from 
some  source.  This  source  must  be  uranium,  for  every  known  ura- 
nium ore  contains  radium  (McCoy)  and  radium  emanation  (Boltwood) 
in  amounts  proportional  to  the  uranium  content.  Boltwood  has  now 
shown  that  radium  is  not  a  direct  product  of  the  decay  of  uranium, 
but  is  formed  from  a  very  slowly  changing  element,  ionium,  which,  in 
turn,  arises  out  of  U  — X. 

It  has  recently  been  found  (Ramsay  and  Soddy)  that  helium  (p.  430) 
is  one  of  the  decomposition  products  of  the  radium  emanation. 

The  phenomena  of  radio-active  substances  lead  undeniably  to  the 
startling  conclusion  that  some,  if  not  all,  of  the  elements  are  capable 
of  spontaneous  decomposition.  The  transmutation  of  the  elements,  in 
the  manner  indicated,  seems  to  be  clearly  established,  since  it  appears 
certain  that  the  element  uranium  produces  another  undoubted  element, 
radium,  which,  in  turn,  yields  the  element  helium.  Besides  the  U  —  Ra 
series  of  radio-active  substances,  others  are  known.  Thus,  all  of  the 
compounds  of  thorium  are  radio-active,  and  other  less  well-character- 
ized substances,  such  as  polonium,  actinium,  radio-tellurium,  radio-lead, 
etc.,  have  been  separated  from  pitchblende.  It  is  probable  that  radio- 
lead  is  radium-D,  and  that  radio-tellurium  is  radium-F.  There  is  also 
considerable  evidence  that  many  other  elements  are  very  slightly  radio- 
active, perhaps  TTJ^  as  active  as  uranium.  This  may  mean  that  all 
elements  undergo  an  extremely  slow,  spontaneous  decay.  At  present 
there  is  no  known  means  of  hastening  the  rate  of  radio-active  change. 

Exercises.  —  1.  Construct  equations  showing  the  interactions  of 
(a]  chromic  oxide  and  aluminium,  (b)  strontium  nitrate  and  potassium 
dichromate  in  solution,  (c)  potassium  hydroxide  and  chromic  hydroxide, 
and  the  reversal  on  boiling,  (d)  chlorine  and  potassium  chromite  in 
excess  of  alkali  (what  is  the  actual  oxidizing  agent  ?). 

2.  What  volume  of  oxygen  at  0°  and  760  mm.  (a)  is  obtainable 
from  one  formula-weight  of  potassium  dichromate  (par.  4,  p.  726),  (b) 
is  required  to  oxidize  one  formula- weight  of  chromous  chloride  ? 

3.  To  what  classes  of  actions  should  you  assign  the  three  methods 
of  making  chromic  oxide  (p.  730)  ? 


CHAPTER   XLIII 
MANGANESE 

The  Chemical  Relations  of  the  Element.  —  Manganese  stands, 
at  present,  alone  on  the  left  side  of  the  eighth  column  of  the  periodic 
table.  The  right  side  is  occupied  by  the  halogens.  It  is  never  uni- 
valent,  as  the  halogens  are,  but  its  heptoxide  Mn207  and  the  corre- 
sponding acid,  permanganic  acid  HMn04,  are  in  many  ways  closely 
related  to  the  heptoxide  of  chlorine  and  perchloric  acid  HC104.  Of 
the  lower  oxides  of  manganese,  MnO  is  basic,  and  Mn20s  feebly  basic. 
Mn02  is  feebly  acidic,  Mn03  more  strongly  so,  and  permanganic  acid 
(from  Mn207)  is  a  very  active  acid.  Contrary  to  the  habit  of  feebly 
acidic  and  feebly  basic  oxides,  such  as  those  of  zinc,  aluminium,  and 
tin,  the  basic  oxides  of  manganese  are  not  at  all  acidic,  and  the  acidic 
oxides  (with  the  possible  exception  of  Mn203)  are  not  also  basic.  There 
are  thus  the  five  following,  rather  well-defined  sets  of  compounds, 
showing  five  different  valences  of  the  element.  Of  these  the  first, 
fourth,  and  fifth  are  the  most  stable  and  the  most  important. 

1.  Manganous  compounds,  MnO,  Mn(OH)2,  MnSO4,  etc.      These 
compounds  resemble  those  of  the  magnesium  family  (and  those  of 
Fe*%).     The  salts  of  weak  acids,  such  as  the  carbonate  and  sulphide, 
are  easily  made,  and  there  is  little  hydrolysis  of  the  halides.     The 
salts  are  pale-pink  in  color. 

2.  Manganic  compounds,   Mn203,   Mn(OH)8,  Mn2(S04)8,    [MnCl8]. 
The  salts  resemble  the  chromic  and  aluminium  salts  in  behavior,  but 
are  even  less  stable  than  those  of  quadrivalent  lead.     They  are  com- 
pletely hydrolyzed  by  little  water.     The  salts  are  violet  in  color. 

3.  Manganites,  Mn02,  H2Mn03,  CaMn08.     The  alkali  manganites 
are  strongly  hydrolyzed,  like  plumbates^and  stannates. 

4.  Manganates,  MnO3,  H2Mn04,  K2Mn04.     The  salts  resemble  the 
sulphates  and  chromates,  but  are  much  more  easily  hydrolyzed.     The 
free  acid  resembles  chloric  acid  in  that  when  it  decomposes  it  yields 
a  higher  acid  (HMn04)  and  a  lower  oxide  (Mn02).     The  salts  are 
green  in  color. 

6.  Permanganates,   Mn207,  HMn04(hydrated),  KMn04.     The  salts 

737 


738  INORGANIC   CHEMISTRY 

resemble  the  perchlorates.  and  are  not  hydrolyzed  by  water.  They  are 
reddish-purple  in  color. 

It  will  be  seen  that  the  element  manganese  changes  its  character 
totally  with  change  in  valence,  and  in  each  form  of  combination 
resembles  some  set  of  elements  of  valence  identical  with  that  which  it 
has  itself  assumed. 

Occurrence  and  Isolation.  —  The  chief  ore  is  the  dioxide,  pyro- 
lusite  Mn02,  which  always  contains  compounds  of  iron.  Other  man- 
ganese minerals  are :  braunite  Mn203 ;  the  hydrated  form,  manganite 
MnO(OH) ;  hausmannite  Mn304;  and  manganese  spar  MiiC03.  The 
last  is  isomorphous  with  calcite.  The  metal  is  most  easily  made  by 
reducing  one  of  the  oxides  with  aluminium  by  Groldschmidt's  method. 

Physical  and  Chemical  Properties. — The  metal  manganese 
has  a  grayish  luster  faintly  tinged  with  red.  It  rusts  in  moist  air,  and 
easily  displaces  hydrogen  from  dilute  acids,  giving  manganous  salts. 
Its  alloys  with  iron,  such  as  ferro-manganese  (20-80  per  cent  mangan- 
ese), are  used  in  the  arts. 

Oxides.  —  Manganous  oxide  MnO  is  a  green  powder,  made  by  re- 
ducing any  of  the  other  oxides  with  hydrogen.  Hausmannite  Mn304 
is  red.  An  oxide  having  this  composition  is  formed  when  any  of  the 
other  oxides  is  heated  in  air,  oxidation  or  reduction,  as  the  case  may 
be,  taking  place  (cf.  p.  701).  This  oxide  corresponds  to  minium 
Pb804  (p.  701)  rather  than  to  Fe8O4,  for  with  dilute  acids  it  gives  a 
soluble  manganous  salt  and  a  precipitate  of  the  dioxide : 

Mn2MnO4  +  4HN08  ->  2Mn(N08)2  +  H4MnO4  J. 

The  hydrated  dioxide  H4Mn04  subsequently  loses  water.  Haus- 
mannite also  forms  square  prismatic  crystals.  In  view  of  its  behavior 
with  acids  and  its  crystalline  form,  it  is  thought  to  be  an  orthoman- 
ganite  of  manganese  Mn2Mn04,  rather  than  a  derivative  of  manganic 
oxide,  Mn(Mn02)2,  which  would  be  a  spinelle  (p.  686).  The  magnetic 
oxide  of  iron  Fe  (Fe02)2  belongs  to  the  regular  system,  like  the  spinelles. 
Manganic  oxide  Mn208  is  brownish-black,  and  is  formed  by  heating  any 
of  the  oxides  in  oxygen.  In  dilute  acids  it  behaves  as  if  it  were  a 
manganite  of  manganese  Mn.Mn03,  for  it  gives  a  manganous  salt  and 
manganese  dioxide.  Yet  compounds  of  trivalent  manganese  are  known, 
and  this  may  be  one. 


MANGANESE  739 

Manganese  dioxide  Mn02  is  black,  and  is  most  easily  prepared  in 
pure  condition  by  gentle  ignition  of  manganous  nitrate.  The  hydrated 
forms  of  the  oxide  are  produced  by  reactions  like  those  just  mentioned, 
and  by  adding  a  hypochlorite  or  hypobromite  to  manganous  hydroxide 
suspended  in  water.  Manganese  dioxide  is  not  a  peroxide  in  the  re- 
stricted sense  (cf.  p.  308).  It  is  used  for  manufacturing  chlorine, 
although  electrolytic  processes  are  now  driving'  it  out  of  this  field.  In 
glass-making  (q.v.),  it  is  employed  to  oxidize  the  green  ferrous  silicate, 
.derived  from  impurities  in  the  sand,  to  the  pale-yellow  ferric  com- 
pound. The  amethyst  color  of  the  manganic  silicate  which  is  formed 
tends  to  neutralize  this  yellow.  The  dioxide  forms  the  depolarizer  in 
the  Leclanche  cell  (p.  673). 

Manganese  trioxide  is  a  red,  unstable  powder.  Manganese  hept- 
oxide  is  a  brownish-green  oil  (see  below). 

When  any  of  these  oxides  is  heated  with  an  acid,  a  manganous  salt 
is  obtained.  Salts  of  this  class  are,  in  fact,  the  only  stable  substances 
in  which  manganese  is  combined  with  an  acid  radical.  In  this  action 
the  oxides  containing  more  oxygen  than  does  MnO  give  off  oxygen,  or 
oxidize  the  acid  (cf.  p.  172).  When  the  oxides  are  heated  with  bases, 
in  the  presence  of  air,  manganates  are  always  formed.  With  the 
oxides  containing  a  smaller  proportion  of  oxygen  than  Mn08  oxygen 
is  taken  from  the  air. 

Manganous  Compounds.  —  The  manganous  salts  are  formed  by 
the  action  of  acids  upon  the  carbonate  or  any  of  the  oxides.  Thus  the 
chloride  MnCl2, 4H20  is  obtained  in  pale-pink  crystals  from  a  solution 
made  by  treating  the  dioxide  with  hydrochloric  acid  and  driving  off 
the  chlorine  liberated  by  oxidation  (p.  171).  The  hydroxide  Mn(OH)2 
is  formed  as  a  white  precipitate  when  a  soluble  base  is  added  to  a  solu- 
tion of  a  manganous  salt.  This  body  passes  into  solution  when  ammo- 
nium salts  are  added,  and  cannot  be  precipitated  in  their  presence  on 
account  of  the  formation  of  molecular  ammonium  hydroxide  and  the 
suppression  of  hydroxidion  (cf.  magnesium  hydroxide,  p.  644).  The 
hydroxide  quickly  darkens  when  exposed  to  the  air  and  passes  over 
into  hydrated  manganic  oxide  MnO  (OH). 

Manganous  sulphate  gives  pink  crystals  of  a  hydrate.  Below  6° 
the  solution  deposits  MnS04, 7H20,  which  is  a  vitriol  (p.  649).  Be- 
tween 7°  arid  20°  the  product  is  MnS04, 5H20,  asymmetric  and  isomor- 
phous  with  CuS04, 5H20.  Above  25°  monosymmetric  prisms  of 
MnS04, 41^0  are  obtained-  These  hydrates  have  different  aqueous 


740  INORGANIC   CHEMISTRY 

tensions  and  may  be  formed  from  one  another  by  lowering  or  raising 
the  pressure  of  water  vapor  around  the  substance  (p.  122).  The  signi- 
ficance of  the  temperatures  proper  to  the  crystallization  of  each  (cf.  pp. 
573,  608,  623)  is  that  a  given  solid  hydrate  can  be  formed  only  in  a  solu- 
tion which  is  saturated  with  respect  to  that  hydrate  and  has  the  same 
aqueous  tension  as  the  hydrate.  These  conditions  are  necessary  to  that 
state  of  equilibrium  between  the  solution  and  the  hydrate  on  which  the 
co-existence  of  solution  and  hydrate  during  crystallization  depends  (cf. 
p.  160).  Hence  the  hydrates  with  the  larger  proportions  of  water,  and 
the  higher  aqueous  tensions,  are  formed  in  the  colder  solutions  which 
contain  less  cf  the  solute  when  saturated  and  have  therefore  at  a  given 
temperature  themselves  relatively  high  aqueous  tensions. 

The  presence  of  a  foreign  dissolved  body,  since  it  will  lower  the 
'  vapor  tension  of  the  solution,  may  similarly  cause  the  formation  of  a 
lower  hydrate.  Thus,  at  the  ordinary  temperature,  calcium  sulphate  solu- 
tion has  a  higher  aqueous  tension  than  gypsum,  and  therefore  gypsum 
is  deposited  from  it,  and  anhydrite  will  turn  into  gypsum  if  placed  in  it. 
But  calcium  sulphate  solution  containing  much  of  the  chlorides  of 
sodium  and  magnesium  has  a  lower  aqueous  tension  than  gypsum,  and 
so  anhydrite  is  deposited,  and  gypsum  in  contact  with  such  a  solution 
would  lose  its  water  of  hydration.  This  explains  the  deposition  of 
anhydrite  in  the  salt  layers  (cf.  p.  603). 

Manganous  carbonate  MnC03  is  a  white  powder  formed  by  pre- 
cipitation. The  sulphide  MnS  is  obtained  as  a  green  powder  by  lead- 
ing hydrogen  sulphide  over  any  of  the  oxides.  A  flesh-colored, 
hydrated  form  MnS,  H20  is  more  familiar  and  is  precipitated  by  am- 
monium sulphide  from  manganous  salts.  It  interacts  with  mineral 
acids  and  even  with  acetic  acid,  so  that  it  cannot  be  precipitated  by 
hydrogen  sulphide  (cf.  p.  651). 

The  manganous  salts  of  weak  acids,  such  as  the  carbonate  and 
sulphide,  darken  when  exposed  to  air  and  are  oxidized,  with  formation 
of  hydrated  manganic  oxide.  As  we  have  seen,  manganous  hydroxide 
is  similarly  oxidized  and  these  salts  are  precisely  the  ones  which 
/should  furnish  the  hydroxide  by  hydrolysis.  While  there  is  a  general 
resemblance  between  the  manganous  salts  and  the  stannous,  chromous, 
and  ferrous  salts,  the  manganous  salts  of  active  acids  are  not  oxidized 
by  the  air  as  are  the  corresponding  salts  of  the  other  three  metals. 

Manganic  Compounds*  —  The  base  of  this  set  of  compounds, 
manganic  hydroxide  Mn  (OH)8,  is  slowly  deposited  by  the  action  of 


MANGANESE  741 

the  air  on  an  ammoniacal  solution  of  a  manganous  salt  in  salts  of 
ammonium.  The  chloride  MnCl8  is  present  in  the  liquid  obtained  by 
the  action  of  hydrochloric  acid  upon  manganese  dioxide  (cf.  p.  171),  but 
loses  chlorine  very  readily  and  cannot  be  isolated.  Double  salts  such  as 
MnCl3,  2KC1  and  MnF3, 2KF,  2H2O  are  known.  Manganic  sulphate 
Mn2(S04)3  is  deposited  as  a  violet-red  powder  when  hydrated  man- 
ganese dioxide  is  heated  with  concentrated  sulphuric  acid  at  160°. 
It  is  deliquescent  and  is  rapidly  hydrolyzed  in  the  cold  even  by  a 
little  water,  giving  the  brownish-black  hydroxide : 

Mn2(S04)3  -f  6H2O  ->  2Mn(OH)3  +  3H2SO4. 

The  caesium-manganic  alum  Cs2S04,Mn2(S04)8,  24H2O  seems  to  be  the 
most  stable  derivative. 

Manganites,  — Although  manganese  dioxide  interacts  when  fused 
with  potassium  hydroxide,  simple  salts  derived  from  H2Mn08  ( =  H20, 
Mn02)  or  H4Mn04  ( =  2H20,Mn02)  are  not  formed.  The  products  are 
complex,  as  K2Mn5On.  Some  less  complex  manganites  are  formed  in 
the  Weldon  process  for  utilizing  the  manganous  chloride  obtained  in 
manufacturing  chlorine.  The  liquor  is  mixed  with  slaked  lime,  and 
air  is  blown  through  the  mass  of  calcium  and  manganous  hydroxides 
which  is  thus  obtained.  Black  manganites  of  calcium,  such  as 
CaMn03(=  CaO,Mn02)  and  CaMn206(CaO,2Mn02)  are  thus  formed: 

Ca(OH)2  +  2Mn(OH)2  +  02  ->  CaMn205  +  3H20, 

and  when  afterwards  treated  with  hydrochloric  acid  they  behave  like 
mixtures  of  manganese  dioxide  and  calcium  oxide.  As  we  have  seen 
(p.  738),  the  oxides  Mn304  and  Mn208  may  be  manganites  of  manga- 
nese. 

Manganates.  —  When  one  of  the  oxides  of  manganese  is  fused 
with  potassium  carbonate  arid  potassium  nitrate  a  green  mass  is 
obtained.  The  green  aqueous  extract  deposits  potassium  manganate 
K2Mn04  in  rhombic  crystals,  which  are  isomorphous  with  those  of  potas- 
sium sulphate,  and  are  almost  black : 

K2C03  +  Mn02  -f  0  -+  K2Mn04  +  C02. 

The  acid  H2Mn04,  itself  unknown,  must  be  weak,  for  the  potassium 
salt  is  easily  hydrolyzed.  The  salt  remains  unchanged  in  solution 
only  in  presence  of  free  alkali,  the  hydroxidion  of  the  alkali  combin- 


742  INORGANIC   CHEMISTRY 

ing  with  and  suppressing  the  hydrion  of  the  water  whose  combination 
with  the  Mn04"  ion  constitutes  the  hydrolysis.  When  the  concentra- 
tion of  the  hydroxidion  is  reduced  by  dilution,  or,  better  still,  when  a 
weak  acid  such  as  carbonic  acid  or  acetic  acid  is  used  to  neutralize  it, 
the  salt  is  hydrolyzed,  according  to  the  partial  equation  : 

K2Mn04  4-  2H20  U  2KOH  (  +  H2Mn04).  (1) 

The  free  acid  immediately  changes  so  that  a  part  is  oxidized  to  per- 
manganic acid,  giving  a  purple-red  color  to  the  solution,  and  a  part  is 
reduced  to  manganese  dioxide,  giving  a  black  precipitate.  The  trans- 
formation is  similar  to  that  of  chloric  acid  (p.  275).  The  equation 
may  be  made  by  noting  that  manganic  acid  has  the  composition  H20, 
Mn08  and  changes  so  as  to  yield  H20,Mn207  and  Mn02.  Thus  each 
molecule  of  H2Mn04,  in  forming  a  molecule  of  Mn02,  yields  one  unit 
of  oxygen,  while  2(H20,Mn03)  4  0  are  required  to  give  H^Mi^O,  4- 
HaO: 

3(H20,Mn08)  ->  H20,  Mn207  +  MnO2  +  2H2O 

or  (3H2Mn04)  ->  2HMn04  +  MnO2  4-  2H20.  (2) 

In  consequence  of  the  presence  of  potassium  hydroxide  (equation  (1)) 
the  product  is  potassium  permanganate  : 

2KOH  +  2HMn04  -»  2KMn04  +  2H20.  (3) 

Multiplying  equation  (1)  by  3,  omitting  the  manganic  acid,  and  add- 
ing the  three  partial  equations,  we  have  the  equation  for  the  action  as 
it  really  occurs  : 


n04  4-  2H20  -i  4KOH  +  2KMn04  +  MnOa. 
In  terms  of  the  ions  the  equation  is  simpler  : 

3Mn04"  +  2H*  ->  20H'  4-  2MnO/  4-  Mn02. 

The  alkaline  solution  of  potassium  manganate  interacts  readily 
with  oxidizable  substances.  Thus  oxalic  acid  is  converted  into  car- 
bonic acid,  and  alcohol  into  acetic  acid.  The  details  of  the  change 
depend  upon  the  amount  of  free  alkali  present  and  the  nature  of  the 
product  of  oxidation.  Lower  oxides  of  manganese  such  as  Mn02  are 
usually  precipitated. 

Permanganates.  —  Potassium  permanganate  KMn04  is  made  by 
hydrolysis  of  the  manganate  as  shown  above,  and  is  obtained,  in  purple 


MANGANESE  743 

crystals  with  a  greenish  luster,  by  evaporation  of  the  solution.  The 
crystals  are  rhombic  prisms,  isomorphous  with  potassium  perchlorate. 
To  avoid  the  loss  of  manganese  thrown  down  as  dioxide,  the  action  is 
carried  out  commercially  by  passing  ozone  through  the  solution  of  the 
manganate  : 

2K2Mn04  +  08  +  H20  ->  2KMn04  +  0,  +  2KOH. 

Sodium  permanganate  is  made  in  a  similar  manner.  It  is  not  obtain- 
able in  solid  form,  but  its  solution  is  known  as  "  Condy's  disinfecting 
fluid."  This  liquid  owes  its  properties  to  the  oxidizing  power  of  the 
salt.  Permanganic  acid  is  a  very  active  acid,  that  is,  it  is  highly  ion- 
ized in  aqueous  solution.  A  solid  hydrate  of  the  acid  may  be  secured 
in  reddish-brown  crystals  by  adding  sulphuric  acid  to  a  solution  of 
barium  permanganate  and  allowing  the  filtrate  to  evaporate : 

Ba(Mn04)2  +  H2S04  +  aH20  <=±  BaSOJ  +  2HMn04,icH20. 

This  hydrate  decomposes,  on  being  warmed  to  32°,  and  yields  oxygen 
and  manganese  dioxide.  When  a  very  little  dry,  powdered  potassium 
permanganate  is  moistened  with  concentrated  sulphuric  acid,  brownish- 
green,  oily  drops  of  permanganic  anhydride  (manganese  heptoxide) 
Mn207  are  formed.  This  compound  is  volatile,  giving  a  violet 
vapor,  and  is  apt  to  decompose  explosively  into  oxygen  and  manganese 
dioxide.  Its  oxidizing  power  is  such  that  combustibles  like  paper, 
ether,  and  illuminating-gas  are  set  on  fire  by  contact  with  it. 

Potassium  permanganate  is  much  used  for  oxidations.  The  actions 
are  different  according  as  the  substance  is  employed  (1)  in  alkaline, 
(2)  in  acid,  or  (3)  in  neutral  solution. 

1.  When  an  alkali,  such  as  potassium  hydroxide,  is  added,  the 
action  by  which  the  permanganate  is  formed  is  reversed,  and  the  solu- 
tion becomes  green  from  the  production  of  the  manganate : 

4KMn04  +  4KOH  ->  4K2Mn04  +  2H20  +  02, 
or  4MnO/  +  40H'  ->  4Mn04"  +  2H20  +  O2. 

When  a  substance  capable  of  being  oxidized  is  present,  the  reduction 
proceeds  further  and  manganese  dioxide  is  precipitated.  Schemati- 
cally :  Mn207  — >  2Mn02  +  30,  so  that  two  molecules  of  the  perman- 
ganate, in  alkaline  solution,  can  furnish  three  chemical  units  of 
oxygen  to  the  oxidizable  body. 


744  INORGANIC   CHEMISTRY 

2.    With  an  acid,  the  amount  of  oxygen  available  is  greater,  for  the 
manganous  salt  of  the  acid  is  formed  : 

Mn207  -*  2MnO  +  50. 

Thus,  when  sulphuric  acid  is  added  to  potassium  permanganate  solu- 
tion, and  sulphur  dioxide  is  led  through  the  mixture,  we  have  : 

2KMn04  +  3H2S04  ->  K2S04  +  2MnS04  +  3H20(+  50)    (1) 
S08->5H2S04  (2) 

* 


__ 
2KMn04  +  3H2S04  +  5H2S03-*K2S04  +  2MnS04  +  3H20  -|-5H2S04 

In  this  case,  since  sulphuric  acid  is  a  product,  the  preliminary  addi- 
tion of  the  acid  was  superfluous.  In  other  cases,  the  partial  equation 
(1),  showing  the  available  50,  remains  the  same,  while  the  other  par- 
tial equation  varies  with  the  substance  being  oxidized.  Thus,  with 
hydrogen  sulphide  as  reducing  agent,  we  have  : 

(0)  +  H2S  -»  H20  +  S     x  5  (2') 

and  with  ferrous  sulphate,  we  get  ferric  sulphate  : 

2FeS04  +  H2S04(  +  0)  ->  Fe2(S04)8  +  H20     x  5         (2") 

As  before  (2')  and  (2")  must  be  multiplied  throughout  by  five,  before 
summation  is  made.  Since  the  permanganate  is  deep-purple  in  color, 
while  the  manganous  salt  is  almost  colorless,  this  sort  of  action  can  be 
used  without  an  indicator  for  quantitative  experiments.  The  stan- 
dard solution  of  the  permanganate  is  added  from  a  burette  until  the 
purple  color  ceases  to  disappear  ;  and  the  amount  used  enables  us  to 
calculate  the  quantity  of  ferrous  salt,  oxalic  acid,  nitrous  acid,  or  other 
oxidizable  substance,  which  was  present.  The  last  named  substance 
is  oxidized  to  nitric  acid. 

3.    When  dry  potassium  permanganate  is  heated,  it  decomposes 
as  follows  : 

2KMn04  ->  K2Mn04  +  Mn02  +  O2. 

The  neutral  solution  resembles  that  of  potassium  dichromate  in  oxi- 
dizing substances  which  are  reducing  agents,  but  is  more  active. 
Thus  when  the  powdered  salt  is  moistened  with  glycerine,  the  mass 
presently  bursts  into  flame.  The  fingers  are  stained  brown,  receiving 
a  deposit  of  manganese  dioxide,  in  consequence  of  the  reducing  power 
of  the  unstable  organic  substances  in  the  skin.  The  destruction 
of  minute  organisms  by  Condy's  fluid  results  from  a  similar  action. 


MANGANESE  745 

Analytical  Reactions  of  Manganese  Compounds.  —  The  ions 
commonly  encountered  are  dimanganion  Mn  ,  which  is  very  pale-pink 
in  color,  permangananion  MnO/,  which  is  purple,  and  mangananion 
MnO/',  which  is  green. 

The  manganous  compounds  give  with  ammonium  sulphide  the 
flesh-colored,  hydrated  sulphide  which  is  soluble  in  acids.  Bases  give 
the  white  hydroxide,  which  darkens  by  oxidation,  and  is  soluble  in 
salts  of  ammonium.  The  black,  hydrated  dioxide  is  precipitated  by 
hypochlorites. 

All  compounds  of  manganese  confer  upon  the  borax  bead  an  ame- 
thyst color  which,  in  the  reducing  flame,  disappears.  A  bead  of 
sodium  carbonate  and  niter  becomes  green  on  account  of  the  formation 
of  the  manganate. 

Exercises.  —  1.  Consider  the  valence  of  manganese  in  the  oxides 
Mn304  and  Mn203,  on  the  theory  that  they  are  manganites. 

2.  What  do  we  mean  by  saying  that  (a)  chromous  chloride  is  stable 
(p.  119),  but  easily  oxidized  by  the  air,  (b)  permanganic  acid  is  an 
active  acid,  (c)  permanganic  acid  is  an  active  oxidizing  agent  in  pres- 
ence of  an  acid  ? 

3.  Formulate  the  oxidations  of  oxalic  acid  and  of  nitrous  acid  by 
potassium  permanganate  in  acid  solution. 


CHAPTER  XLIV 

IRON,  COBALT,  NICKEL 

THE  elements  iron  (Fe,  at.  wt.  55.9),  cobalt  (Co,  at.  wt.  59),  and 
nickel  (Ni,  at.  wt.  58.7)  are  not  corresponding  members  of  successive 
periods,  like  the  families  hitherto  considered.  They  are  neighboring 
members  of  the  first  long  period,  lying  between  its  first  and  second 
octaves  (p.  408),  and  form  a  transition  group  between  the  adjoining 
elements  within  those  octaves.  Thus,  iron  forms  ferrates  M9IFeVI04 
and  ferric  salts  FemCl8,  as  well  as  ferrous  salts  Fe11^.  These 
resemble  the  chromates  and  manganates,  the  chromic  and  manganic 
salts,  and  the  chromous  and  manganous  salts,  respectively.  Cobalt 
forms  cobaltic  and  cobaltous  salts,  like  Co2m(S04)3  and  Co11^.  Nickel 
enters  only  into  nickelous  salts,  like  MC12,  and  thus  links  iron  and 
cobalt  with  copper  and  zinc  which  are  both  bivalent  elements.  The 
free  metals  of  this  family  are  magnetic,  iron  showing  this  property 
strongly  and  cobalt  very  distinctly. 

IRON. 

Chemical  Relations  of  the  Element.  —  The  oxides  and  hydrox- 
ides FeO  and  Fe(OH)2,  Fe203  and  Fe(OH)3  are  basic,  the  former  more 
strongly  so  than  the  latter.  The  ferrous  salts,  derived  from  Fe(OH)2, 
resemble  those  of  the  magnesium  group  and  those  of  Cr**  and  Mn**  and 
are  little  hydrolyzed.  The  ferric  salts,  derived  from  Fe(OH)8,  re- 
semble those  of  Cr""  and  Al**"  and  are  hydrolyzed  to  a  considerable 
extent.  Ferric  hydroxide  is  even  less  acidic,  however,  than  is  chromic 
hydroxide.  Iron  gives  also  a  few  ferrates  K2Fe04,  CaFe04,  etc.,  de- 
rived from  an  acid  H2FeO4  which,  like  manganic  acid  H2Mn04  (p.  742), 
is  too  unstable  to  be  isolated.  Complex  anions  containing  this  element, 
such  as  the  anion  of  K4.Fe(CN)6,  are  familiar,  but  complex  cations 
containing  ammonia  are  unknown. 

The  ferrous  salts  differ  from  most  of  the  manganous  salts  and  re- 
semble the  chromous  and  stannous  salts  in  being  easily  (although  not 
quite  so  easily)  oxidized  by  the  air.  They  pass  into  the  ferric 
condition. 

746 


IRON,    COBALT,   NICKEL 


747 


Occurrence.  —  Free  iron  is  found  in  minute  particles  in  some 
basalts,  and  many  meteorites  are  composed  of  it.  Meteoric  iron  can 
be  distinguished  from  specimens  of  terrestrial  origin  by  the  fact  that 
it  contains  3-8  per  cent  of  nickel.  The  chief  ores  of  iron  are  the 
oxides,  haematite  Fe2O8  and  magnetite  Fe804,  and  the  carbonate  FeC03, 
siderite.  The  first  is  reddish  and  columnar  in  structure ;  but  black, 
shining,  rhombohedral  crystals,  known  as  specularite,  are  also  found. 
Hydrated  forms,  like  brown  iron  ore  2Fe208,  3H20,  are  also  common. 
Siderite  is  pale-brown  in  color  and  rhombohedral,  isomorphous  with 
calcite.  When  mixed  with  clay  it  forms  iron-stone  and,  with  20-25 
per  cent  of  coal  in  addition,  black-band.  Pyrite  FeS2  consists  of 
golden-yellow,  shining  cubes  or  pentagonal  dodecahedra.  It  is  used, 
on  account  of  its  sulphur,  in  the  manufacture  of  sulphuric  acid,  but, 
from  the  oxidized  residue,  iron  of  sufficient  purity  is  obtained  with 
difficulty.  Compounds  of  iron  are  con- 
tained in  the  haemoglobin  of  the  blood, 
and  doubtless  plays  an  important  part 
in  connection  with  the  vital  func- 
tions of  this  substance.  By  interaction 
with  organic  compounds  of  iron  present 
in  the  tissues,  ammonium  sulphide 
blackens  the  skin,  ferrous  sulphide 
being  formed. 

Metallurgy.  —  The  ores  of  iron  are 
usually  first  roasted  in  order  to  decom- 
pose carbonates  and  oxidize  sulphides. 
They  are  then  reduced  with  coke.  Ores 
containing  lime  or  magnesia  are  mixed 
with  an  acid  flux,  such  as  sand  or  clay- 
slate,  in  order  that  a  fusible  slag  may 
be  formed.  Conversely,  ores  containing 
silica  and  clay  are  mixed  with  lime- 
stone. With  proper  adjustment  of  the 
ingredients  the  process  can  be  carried 
on  continuously  in  a  blast  furnace  (Fig.  FIG.  105. 

105).      The  solid  materials  thrown   in 

at  the  top  are  converted,  as  they  slowly  descend,  completely  into 
gases  which  escape  and  liquids  (iron  and  slag)  which  are  tapped 
off  at  the  bottom.  Heated  air  is  blown  in  at  the  bottom  through 


748  INORGANIC   CHEMISTRY 

tuyeres,  and  the  top  is  closed  by  a  bell  which  descends  for  a  moment 
when  an  addition  is  made  to  the  charge.  The  gases,  which  contain 
much  carbon  monoxide,  are  led  off  and  used  to  heat  the  blast  or  to 
drive  gas-engines. 

The  main  action  takes  place  between  the  carbon  monoxide,  present 
in  consequence  of  the  excess  of  carbon,  and  the  oxide  of  iron : 

Fe804  +  4CO  «=»  3Fe  -f  4C02. 

Since  the  action  is  a  reversible  one,  a  large  excess  of  carbon  monoxide 
is  required.  Different  smelters  use  different  proportions,  the  ratio  of 
the  amount  actually  used  to  that  supplied  varying  from  1 : 2  to  1 :  15. 
The  actual  ratio  by  volume  of  carbon  dioxide  to  carbon  monoxide 
required  for  equilibrium  with  the  two  solids  is,  61 :  39  at  650°,  and 
7 : 93  at  800°. 

In  the  upper  part  of  the  furnace,  the  heat  (400°)  loosens  the  texture 
of  the  ore.  Further  down,  the  temperature  is  higher  (500-900°), 
and  the  carbon  monoxide  reduces  the  oxide  of  iron  to  particles  of  soft 
iron.  A  temperature  high  enough  to  melt  pure  iron  is  barely  reached 
anywhere  in  the  furnace,  but,  a  little  lower  down,  by  union  with  carbon, 
the  more  fusible  cast  iron  (1200°)  is  formed  and  falls  in  drops  to  the 
bottom.  It  is  in  this  region  also  that  the  slag  is  produced.  If  the  flux 
had  begun  sooner  to  interact  with  the  unreduced  ore,  iron  would  have 
been  lost  by  the  formation  of  the  silicate.  The  iron  collects  below 
the  slag,  and  the  latter  flows  continuously  from  a  small  hole.  The 
former  is  tapped  off  at  intervals  of  six  hours  or  so  from  a  lower 
opening. 

Cast  Iron  and  Wrought  Iron.  —  Pure  iron  is  not  manufactured, 
and  indeed  would  be  too  soft  for  most  purposes.  Piano-wire,  however,, 
is  about  99.7  per  cent  pure.  The  product  obtained  from  the  blast  fur- 
nace contains  92-94  per  cent  of  iron  along  with  2.6-4.3  per  cent  of  car- 
bon, often  nearly  as  much  silicon,  varying  proportions  of  manganese, 
and  some  phosphorus  and  sulphur.  The  last  four  ingredients  are  lib- 
erated from  combination  with  oxygen  by  the  carbon  in  the  hottest  part 
of  the  furnace  and  combine  or  alloy  themselves  with  the  iron.  Cast 
iron  does  not  soften  before  melting,  as  does  the  purer  wrought  iron,  but 
melts  sharply  at  1150-1250°  according  to  the  amount  of  foreign  material 
it  contains.  When  suddenly  cooled  it  gives  chilled  cast  iron  which  is 
very  brittle  and  looks  homogeneous  to  the  eye,  all  the  carbon  being 
present  in  the  form  of  carbide  of  iron  Fe8C  in  solid  solution  in  the 


IRON,    COBALT,    NICKEL 


749 


metal.  By  slower  cooling,  time  is  permitted  for  the  separation  of  part 
of  the  carbon  as  graphite  and  for  other  changes  (see  below),  and  gray 
cast  iron  results.  Spiegel  iron  is  cast  iron  made  from  ores  containing 
5-20  per  cent  of  manganese  and  the  usual  proportion  of  carbon.  Ferro- 
manganese  contains  20-80  per  cent  of  the  same  element. 

Wrought  iron  is  made  by  heating  the  broken  "  pigs  "  of  cast  iron 
upon  a  layer  of  material  containing  oxide  of  iron  and  hammer-slag 
(basic  silicate  of  iron)  spread  on  the  bed  of  a  reverberatory  furnace. 
The  carbon,  silicon,  and  phosphorus  combine  with  the  oxygen  of  the 
oxide,  and  the  last  two  pass  into  the  slag.  The  sulphur  is  found  in  the 
slag  as  ferrous  sulphide.  On  account  of  the  effervescence  due  to 
the  escape  of  carbon  monoxide,  the 
process  is  called  "  pig-boiling. "  The 
iron  is  stirred  with  iron  rods  ("  pud- 
dled") and  stiffens  as  it  becomes 
purer,  until  finally  it  can  be  with- 
drawn in  balls  ("  blooms  ")  and  freed 
from  slag  under  the  steam-hammer. 
It  now  softens  sufficiently  for  weld- 
ing below  1000°  and  melts  at  1550° 
or  lower,  according  to  its  purity. 
If  it  still  contains  more  than  a  trace 

of  combined  phosphorus  it  is  brittle  when  cold  ("  cold  short ").  A 
little  surviving  sulphide  of  iron  makes  it  brittle  when  hot  ("red- 
short  ")  and  unsuitable  for  forging.  Wrought  iron  should  contain 
only  0.1-0.2  per  cent  of  carbon.  The  above  operations  are  now  largely 
performed  by  machinery. 

Steel.  —  This  is  a  variety  of  iron  almost  free  from  phosphorus, 
sulphur,  and  silicon.  Tool  steel  contains  0.9-1.5  per  cent  of  carbon, 
structural  steel  0.2-0.6  per  cent,  and  mild  steel  0.2  per  cent  or  even 
less.  Steel  combines  the  properties  of  cast  and  wrought  iron,  being 
hard  and  elastic,  and  at  the  same  time  available  for  forging  and  weld- 
ing when  the  proportion  of  carbon  is  low. 

Steel  is  made  largely  by  the  Bessemer  process.  The  molten  cast 
iron  is  poured  into  a  converter  (Fig.  106)  and  a  blast  of  air  (A)  is 
blown  through  it.  The  oxidation  of  the  manganese,  carbon,  silicon, 
and  perhaps  a  little  of  the  iron  gives  out  sufficient  heat  to  raise  the 
temperature  of  the  mass  above  the  melting-point  of  wrought  iron. 
The  required  proportion  of  carbon  is  then  introduced  by  adding  pure 


FIG.  106. 


750  INORGANIC   CHEMISTRY 

cast  iron,  spiegel  iron,  or  coke,  and  the  contents,  first  the  slag,  and 
then  the  molten  steel,  are  finally  poured  into  molds  by  turning  the 
converter.  When  the  cast  iron  contains  much  phosphorus,  the  oxide 
of  this  element  is  reduced  again  by  the  iron  as  fast  as  it  is  formed  by 
the  blast.  In  such  cases  a  basic  lining  containing  lime  and  magnesia 
takes  the  place  of  the  sand  and  clay  lining  of  the  ordinary  Bessemer 
converter,  and  a  slag  containing  a  basic  phosphate  of  calcium  is  pro- 
duced. This  modification  constitutes  what  is  known  as  the  Thomas- 
Gilchrist  process.  The  slag  ("  Thomas-slag  ")  when  pulverized  forms 
a  valuable  fertilizer  (cf.  p.  605). 

In  the  Siemens-Martin,  or  open  hearth  process,  the  cast  iron  is 
melted  in  a  saucer-shaped  depression  lined  with  sand,  and  scraps  of 
iron  plate  (for  dilution)  and  haematite,  or  some  other  oxide  ore,  are 
then  added  in  proper  proportions.  The  materials  are  heated  with  gas 
fuel  for  8-10  hours  until  a  sample  shows,  under  the  hammer,  that  the 
process  is  complete.  The  product  is  then  drawn  off  through  a  hole 
and  cast  in  molds. 

Properties  of  Steel*  —  When  steel  is  heated  to  redness  and 
cooled  slowly,  it  is  comparatively  soft.  Sudden  chilling,  however,  ren- 
ders it  harder  than  glass.  By  subsequent,  cautious  heating  the  hardness 
may  be  reduced  to  any  required  extent,  and  this  treatment  is  called 
"  tempering."  The  sufficiency  of  the  heating  is  judged  roughly  by  the 
interference  colors  caused  by  the  thin  film  of  oxide  which  forms  on  the 
surface.  Thus  a  pale-yellow  color  (430-450°  F)  serves  for  tempering 
razors,  a  decided  yellow  (470°)  for  pen-knives,  a  brown  (490-510°)  for 
shears,  a  purple  (520°)  for  table-knives,  a  blue  (530-570°)  for  watch- 
springs  and  sword-blades,  and  a  black-blue  (610°)  for  saws.  Except 
in  the  case  of  watch-springs,  these  films  are  afterwards  removed  by  the 
grinding. 

To  understand  this  behavior  it  must  be  noted  that  there  are  three 
states  of  solid  iron  resembling  the  rhombic  and  monoclinic  states  of 
sulphur  (cf.  p.  368).  The  form  stable  below  765°  is  known  as  a-ferrite 
(wrought  iron).  It  is  magnetic  and  can  hold  little  carbide  of  iron  in 
solid  solution.  Above  765°  this  changes  into  /3-ferrite  which,  likewise, 
holds  little  of  the  carbide  in  solution,  but  is  not  magnetic.  At  890° 
this  changes  into  y-ferrite,  a  non-magnetic  form  in  which  the  carbide  is 
soluble.  When  allowed  to  cool,  iron  assumes  these  forms  in  the  reverse 
order.  If,  now,  a  fluid  solution  of  carbon  in  iron,  suitable  for  steel,  is 
suddenly  chilled,  a  great  part  of  the  cold  mass  is  a  supercooled  solid 


IRON,   COBALT,    NICKEL  751 

solution  of  carbon  in  y-ferride.  This  solid  solution  is  called  marten- 
site  and  is  very  hard  and  brittle.  It  is  less  stable  at  ordinary  tem- 
peratures than  is  a-ferrite,  but,  as  is  the  case  with  yellow  phosphorus 
(p.  359)  and  amorphous  sulphur  (p.  370),  the  low  temperature  having 
once  been  reached,  transformation  into  the  more  stable  form  is  there- 
after exceedingly  slow.  The  material  is  hard  steel. 

When  the  molten  steel  (solution  of  carbon  in  iron)  is  allowed  to 
cool  so  slowly  that  equilibrium  can  be  reached  at  every  step,  a  compli- 
cated series  of  changes  ensues.  First  the  mass  solidifies  (at  or  before 
1130°)  to  a  mixture  of  martensite  (y-ferrite  with  carbon  in  solid  solu- 
tion up  to  2  per  cent)  and  graphite.*  As  the  temperature  now  falls 
very  slowly,  more  graphite  separates  until,  at  1000°,  1.8  per  cent  re- 
mains in  solution.  From  this  point  the  dissolved  carbide  of  iron 
(cementite  Fe3C  containing  6.6  per  cent  of  carbon)  is  separated.  At 
670°  pure  a-ferrite  also  begins  to  appear.  The  final  result  is  a  mechani- 
cal mixture  of  a-ferrite  (wrought  iron),  carbide  of  iron,  and,  if  the 
original  amount  of  carbon  was  sufficiently  large,  graphite.  These  com- 
ponents may  be  recognized  by  making  a  microscopic  study  of  a  polished 
surface,  and  their  formation  may  be  followed  by  chilling  the  specimen 
at  any  desired  stage.  The  soft  iron  which  predominates  in  the  product 
of  slow  cooling  makes  the  whole  soft.  Heating  to  a  high  temperature 
and  sudden  chilling  gives  the  homogeneous  solid  solution  of  carbon 
in  y-ferrite  once  more  and  restores  the  qualities  characteristic  of  steel. 
Moderated  reheating  (tempering)  of  the  chilled  mass  results  in  more 
or  less  partial  accomplishment  of  the  changes  proper  to  slow  cooling, 
and  consequently  in  a  more  or  less  close  approach  to  the  condition 
which  results  from  this. 

The  difference  between  the  effect  of  rapid  and  slow  cooling  of  cast 
iron  (p.  748)  can  now  be  made  clear.  Rapid  cooling  leads  to  the 
omission  of  the  intervening  steps  enumerated  above  and,  if  something 
like  5  or  6  per  cent  of  carbon  is  present,  the  material  turns  almost 
completely  into  the  carbide  (cementite).  This  is  chilled  cast  iron. 
With  slower  cooling,  much  graphite  separates,  and  the  product,  gray 
cast  iron,  contains  much  less  of  the  carbide  and  much  more  free  iron. 

The  various  changes  which  occur  in  cooling  steel  are  retarded  by 
the  presence  of  foreign  substances,  just  as,  with  sulphur  (p.  370), 
foreign  substances  delay  the  change  from  S^  to  S\  and  permit  the 
supercooling  of  the  former  and  its  appearance  in  the  form  of  amorphous 

*  When  molten  cast  iron,  containing  3-4.6  per  cent  of  carbon,  is  cooled  in  this 
fashion,  the  amount  of  graphite  may  be  considerable. 


752  INORGANIC   CHEMISTRY 

sulphur.  Manganese,  nickel,  and  other  metals,  in  particular,  greatly 
reduce  the  facility  with  which  y-ferrite  passes  into  /?-  and  o-ferrite  at 
890°  and  765°.  Thus  iron  with  12  per  cent  of  manganese,  when 
chilled  from  a  high  temperature,  contains  only  supercooled  y-ferrite 
and  is  non-magnetic.  It  has  to  be  kept  for  hours  (instead  of  a  few 
minutes)  at  a  temperature  below  765°,  say  500-600°,  before  it  goes 
over  into  a-ferrite.  Manganese  is  thus  a  valuable  constituent  of  steel  be- 
cause, by  favoring  the  survival  of  the  y-ferrite  in  which  alone  the  carbon 
is  soluble,  it  permits  the  manufacture  of  a  homogeneous  steel  contain- 
ing an  unusually  large  proportion  of  dissolved  carbon,  and  allows 
slower  cooling  without  loss  of  temper. 

Pure  Iron.  —  The  pure  metal  may  be  made  by  reducing  the 
purified  oxalate  in  a  stream  of  hydrogen  or  by  Goldschmidt's  method 
(p.  540). 

Chemical  Properties.  —  When  exposed  to  moist  air,  iron  receives 
a  loosely  adherent  coating  of  rust  (2Fe203,Fe(OH)3).  There  is  still 
uncertainty  as  to  how  the  product  is  formed.  It  may  result  from 
displacement  of  the  hydrogen  of  carbonic  acid,  the  oxygen  assisting 
(cf.  p.  624),  and  subsequent  hydrolysis  of  the  carbonate  and  oxidation 
of  the  ferrous  hydroxide.  The  fact  that  alkalies  prevent  rusting  favors 
this  view,  for  they  should  diminish  the  amount  of  hydrion.  According 
to  another  theory,  water  and  iron  are  simultaneously  oxidized : 

Fe  +  02  +  H20  -*  FeO  +  H202 

and  the  hydrogen  peroxide  immediately  combines  with  the  ferrous 
oxide : 

2FeO  +  H202  ->  Fe202(OH)2.  (2) 

The  presence  of  the  peroxide  cannot  be  demonstrated  in  this  case, 
perhaps  because  it  is  used  up  very  rapidly.  In  the  rusting  of  zinc, 
however,  it  is  always  found  (cf.  also  p.  636). 

Iron  burns  in  oxygen  and  interacts  with  superheated  steam,  giving 
Fe304.  A  superficial  layer  of  this  oxide  adheres  firmly  and  protects 
the  iron  from  the  action  of  the  air  (Barff's  process  for  prevention  of 
rusting). 

Iron  displaces  hydrogen  easily  from  dilute  acids.  Steel  and  cast 
iron,  which  contain  iron,  its  carbide,  and  graphite,  give  with  cold  dilute 
acids  almost  pure  hydrogen,  and  the  carbide  and  graphite  remain  un- 
attacked.  More  concentrated  acids,  however,  particularly  when  warm, 


IRON,    COBALT,    NICKEL  753 


give  off,  along  with  hydrogen,  hydrocarbons  formed  by  interaction 
with  the  carbide  (p.  543).  The  odor  of  the  gas  is  due  to  com- 
pounds of  sulphur  and  phosphorus.  With  dilute  nitric  acid,  iron  gives 
ferrous  nitrate  and  ammonium  nitrate  (cf.  tin,  p.  694)  and  with  the 
concentrated  nitric  acid  ferric  nitrate  and  oxides  of  nitrogen.  It  has 
little  action  upon  alkalies. 

After  being  dipped  in  very  concentrated  nitric  acid,  iron  becomes 
passive  (cf.  chromium,  p.  723),  and  no  longer  displaces  hydrogen  and 
other  elements  lying  below  it  in  the  electromotive  series.  A  sharp 
blow,  however,  produces  a  change  which  spreads  over  the  surface  from 
the  point  struck,  and  the  metal  becomes  active  once  more. 

Ferrous  Compounds.  —  Ferrous  chloride  is  obtained  as  a  pale- 
green  hydrate  FeCl2,  4H20  by  interaction  of  hydrochloric  acid  with 
the  metal  or  the  carbonate.  The  anhydrous  salt  sublimes  in  colorless 
crystals  when  hydrogen  chloride  is  led  over  the  heated  metal.  At  a 
high  temperature  the  vapor  of  ferrous  chloride  has  a  density  correspond- 
ing to  the  simple  formula  FeCLj,  but  at  lower  temperatures  there  is 
much  association  (p.  242)  and  the  formula  approaches  Fe2Cl4.  In 
solution  the  salt  is  oxidized  by  the  air  to  a  basic  ferric  chloride : 

4Fe"  +  02  +  2H20  **  4Fe"*  +  40 H/ 

In  presence  of  excess  of  the  acid,  normal  ferric  chloride  is  formed. 
With  nitric  acid,  ferric  chloride  and  nitric  oxide  are  produced  (p.  442). 

Ferrous  hydroxide  Fe(OH)2  is  thrown  down  as  a  white  precipitate, 
but  rapidly  becomes  dirty-green  and  finally  brown,  by  oxidation.  It 
dissolves  in  solutions  of  salts  of  ammonium,  being,  like  magnesium 
hydroxide  (p.  644),  sufficiently  soluble  in  water  to  require  an  appreci- 
able concentration  of  OH'  for  its  precipitation.  The  ammonium  salts 
convert  this  into  molecular  ammonium  hydroxide.  Ferrous  oxide  FeO 
is  black,  and  is  formed  by  heating  ferrous  oxalate  in  absence  of  air. 
It  may  be  made  also  by  cautious  reduction  of  ferric  oxide  by  hydrogen 
(at  about  300°),  but  is  easily  reduced  further  to  the  metal. 

Ferrous  carbonate  is  found  in  nature,  and  may  be  made  in  slightly 
hydrolyzed  form  by  precipitation.  The  precipitate  is  white  but  rapidly 
darkens  and  finally  becomes  brown,  the  ferrous  hydroxide  produced  by 
hydrolysis  being  oxidized  to  the  ferric  condition.  The  salt  interacts 
with  water  containing  carbonic  acid  after  the  manner  of  calcium  car- 
bonate (p.  482),  and  hence  is  found  in  solution  in  natural  (chalybeate) 
waters. 


754  INORGANIC  CHEMISTRY 

I 

Ferrous  sulphide  may  be  formed  as  a  black,  metallic-looking  mass 
by  heating  together  the  free  elements.  It  is  produced  by  precipita- 
tion with  ammonium  sulphide,  but  not  with  hydrogen  sulphide.  It 
interacts  readily  with  dilute  acids.  The  precipitated  form  is  slowly 
oxidized  to  ferrous  sulphate  by  the  air. 

Ferrous  sulphate  is  obtained  by  allowing  pyrites  to  oxidize  in  the 
air  and  leaching  the  residue : 

2FeS2  +  702  +  211,0  ->  2FeS04  +  2H2S04. 

The  liquor  is  treated  with  scrap  iron  and  the  neutral  solution  evapo- 
rated until  a  hydrate  FeS04,  7H20,  green  vitriol,  or  "  copperas,"  is 
deposited.  This  substance  forms  green  crystals  belonging  to  the 
monosymmetric  system,  but  gives  also  mixed  crystals  in  which  it  is 
isomorphous  with  the  rhombic  vitriols  (cf.  Magnesium  and  Zinc  sul- 
phates, p.  649).  The  crystals  are  efflorescent,  and  become  also  brown 
from  oxidation  to  a  basic  ferric  sulphate  : 

4FeS04  +  02  +  211,0  >-+  4Fe(OH)SO4. 

With  excess  of  sulphuric  acid  and  air,  or  an  oxidizing  agent,  such  as 
nitric  acid,  ferric  sulphate  is  formed.  The  anhydrous  salt  forms  a 
molecular  compound  with  nitric  oxide,  and  the  solution  becomes  brown 
when  this  gas  is  led  through  it,  an  unstable  complex  ion,  perhaps 
FeNO",  being  produced  (cf.  p.  443).  The  double  salts  of  the  form 
(NH4)2SO4,FeS04,6H2O  (Mohr's  salt)  are  not  efflorescent,  and  in  solid 
form  are  less  readily  oxidized  than  is  ferrous  sulphate.  The  ferrous  sul- 
phate is  used  in  dyeing  and  in  making  ink.  The  extract  of  nut-galls  con- 
tains tannic  acid,  HC14H9O9,  which,  with  ferrous  sulphate,  gives  ferrous 
tannate.  This  salt  is  oxidized  by  the  air  to  the  ferric  condition,  and  the 
ferric  compound  is  a  fine,  black  precipitate  which  can  be  suspended  in 
a  solution  of  gum-arabic.  The  resulting  material  is  ink.  The  black 
streaks  seen  below  nail-heads  in  oak  and  other  woods  are  due  to  the 
formation  of  ferrous  carbonate  and  its  interaction  with  the  tannic  acid 
in  the  wood. 

Ferric  Compounds.  —  By  leading  chlorine  into  a  solution  of 
ferrous  chloride,  and  evaporating  until  the  proper  proportion  of  water 
alone  remains,  a  yellow,  deliquescent  hydrate  of  ferric  chloride,  FeCl3> 
6H20  is  obtained.  When  this  is  heated  still  further,  hydrolysis  takes 
place  and  the  oxide  remains.  When  chlorine  is  passed  over  heated 
iron,  the  anhydrous  salt  sublimes  in  dark  scales,  which  are  red  by 


IRON,   COBALT,   NICKEL  755 

transmitted  light.  At  a  high  temperature  the  formula  of  the  vapor  is 
FeCl8,  but  at  lower  temperatures,  in  consequence  of  association,  the 
density  increases  and  the  formula  approaches  Fe^lg.  In  solution, 
the  salt,  like  other  ferric  salts,  can  be  reduced  to  the  ferrous  condition 
by  boiling  with  iron : 

2Fe%"  +  Fe  -»  3Fe". 

The  same  reduction  is  effected  by  hydrogen  sulphide  and  by  stannous 
chloride  (cf.  Mercuric  chloride,  p.  655)  : 

2Fe'"-f  S"->2Fe"+SJ. 
2Fe"'+  Sn"  ->  2Fe"+  Sn"". 

The  last  action  shows  that  ferrous  salts  are  less  active  reducing  agents 
than  are  the  stannous  salts.  The  ferric  ion  is  almost  colorless,  the 
yellow-brown  color  of  solutions  of  ferric  salts  being  due  to  the  pres- 
ence of  ferric  hydroxide  produced  by  hydrolysis.  The  color  deepens 
when  the  solution  is  heated,  and  fades  again  very  slowly,  by  reversal 
of  the  action,  when  the  cold  solution  is  allowed  to  stand.  On  the 
other  hand,  the  hydrolysis  may  be  reversed  and  the  color  may  be  almost 
destroyed,  particularly  in  the  case  of  the  nitrate,  when  excess  of  the 
acid  is  added  to  the  solution  : 

Fe(NO3)3  +  3H20  *=;  Fe(OH)3  +  3HN08. 

Ferric  iodide  is  reduced  by  the  hydriodic  acid  produced  by  its  own 
hydrolysis,  and  hence  ferrous  iodide  does  not  unite  with  iodine  to  form 
this  compound.  The  case  is  similar  to  that  of  cupric  iodide  (p.  621). 

Ferric  hydroxide,  Fe(OH)8,  appears  as  a  brown  precipitate  when  a 
base  is  added  to  a  ferric  salt.  It  does  not  interact  with  excess  of  the 
alkali.  In  this  form  the  substance  is  a  hydrogele  (p.  523)  and  dries  to 
the  oxide  without  giving  definite  intermediate  hydrated  oxides.  The 
hydrates,  Fe408(OH)6  (brown  iron  ore)  and  Fe20(OH)4  (bog  iron  ore), 
however,  are  found  in  nature.  The  hydroxide  passes  easily  into  colloidal 
solution  in  a  solution  of  ferric  chloride,  and  by  subsequent  dialysis 
through  a  piece  of  parchment  (cf.  p.  523)  the  salt  can  be  separated,  and 
a  pure  aqueous  solution  of  the  hydroxide  obtained.  This  solution  is 
red  in  color,  shows  no  depression  in  the  freezing-point,  and  is  not  an 
electrolyte.  It  deposits  the  hydroxide  as  a  brown  precipitate  when 
ionogens  are  added  to  the  solution. 

Ferric  oxide,  Fe208,  is  sold  as  "  rouge  "  and  "  Venetian  red."  It  is 
made  from  the  ferrous  sulphate  obtained  in  cleaning  iron  ware  which 


756  INORGANIC   CHEMISTRY 

is  to  be  tinned  or  galvanized,  and  in  other  ways  in  the  arts.  The  salt 
is  allowed  to  oxidize,  and  the  ferric  hydroxide,  thrown  down  by  the 
addition  of  lime,  is  calcined.  A  purer  form  is  produced  by  dry  distil- 
lation of  the  basic  ferric  sulphate,  an  operation  which  used  to  be  under- 
taken on  a  large  scale  for  making  Nordhausen  sulphuric  acid  (p.  388). 
This  oxide  is  not  distinctly  acidic,  but  by  fusion  with  more  basic  ox- 
ides, compounds  like  franklinite  Zn(Fe02)2  may  be  formed.  It  is 
rediiced  by  hydrogen,  at  about  300°  to  ferrous  oxide  (which  catches  fire 
spontaneously  in  the  air)  and  at  700-800°  to  metallic  iron. 

Magnetic  oxide  of  iron  Fe304  or  lodestone  is  found  in  nature,  and 
is  formed  by  the  action  of  air  (hammer-scale),  steam,  or  carbon  dioxide 
on  iron.  It  forms  octahedral  crystals  like  the  spinelles  (p.  686),  and  is 
assumed  to  be  Fe(Fe02)2. 

Ferric  sulphide  may  be  made  by  fusing  together  the  free  elements. 
It  is  obtained  by  precipitation  when  soluble  sulphides  are  added  to 
solutions  of  ferric  salts  (Stokes) : 

Fe2(S04)3  +  (NHJjS-^EeA  +  3(NH4)2SO4. 

Formerly  the  precipitate  was  supposed  to  be  a  mixture :  2FeS  +  S. 

Ferric  sulphate  is  formed  by  oxidation  of  ferrous  sulphate,  and  is 
obtained  as  a  white  mass  by  evaporation.  It  gives  alums,  such  as 
(NH4)2S04,  Fe2(S04)8, 24H20,  which  are  almost  colorless  when  pure,  but 
usually  have  a  pale  reddish-violet  tinge. 

Pyrite.  —  The  mineral  pyrite  FeS2  (Fools'  gold)  is  the  sulphide  of 
iron  which  is  most  stable  in  the  air.  It  is  found  in  nature  in  the  form 
of  glittering,  golden-yellow  cubes,  octahedrons,  and  pentagonal  dodeca- 
hedrons. It  is  not  attacked  by  dilute  acids,  but  concentrated  hydro- 
chloric acid  slowly  converts  it  into  ferrous  chloride  and  sulphur.  It 
is  reduced  by  hydrogen  to  ferrous  sulphide. 

Cyanides.  —  When  potassium  cyanide  is  added  to  solutions  of  fer- 
rous or  ferric  salts,  yellowish  precipitates  are  produced,  but  the  simple 
cyanides  cannot  be  obtained  in  pure  form.  These  precipitates  interact 
with  excess  of  the  cyanide  giving  soluble  complex  cyanides  of  the  forms 
4KCK,Fe(CN)2  and  3KCJST,  Fe(CN)3  respectively.  These  are  called 
ferro-  and  ferricyanide  of  potassium,  respectively. 

Ferrocyanide  of  potassium,  "  yellow  prussiate  of  potash,"  is  made 
by  heating  nitrogenous  animal  refuse,  such  as  blood,  with  iron  filings 


IRON,    COBALT,    NICKEL  757 

and  potassium  carbonate.  The  resulting  mass  contains  potassium 
cyanide  and  ferrous  sulphide,  and  when  it  is  treated  with  warm  water 
these  interact  and  produce  the  ferrocyanide  : 

2KCN  +  FeS  -4  Fe(CN)2  +  K2S, 
4KCN  +  Fe(CN)2  ->  K4.Fe(CN)6. 

The  salt  is  made  also  from  the  cyanogen  contained  in  crude  illumi- 
nating-gas. It  forms  large,  yellow,  monosymmetric  tables  with  three 
molecules  of  water  of  hydration.  The  solution  contains  almost  ex- 
clusively the  ions  K*  and  Fe(CN)6"",  and  gives  none  of  the  reactions 
of  the  ferrous  ion  Fe".  The  corresponding  acid  H4Fe(CN)6  may  be 
obtained  as  white  crystalline  scales  by  addition  of  an  acid  and  of 
ether  (in  which  the  substance  is  less  soluble  than  in  water)  to  the  salt. 
The  acid  is  a  fairly  active  one  but  is  unstable  and  decomposes  in  a 
complex  manner.  Other  ferrocyanides  may  be  made  by  precipita- 
tion. That  of  copper  Cu2Fe(CN)6  is  brown,  and  ferric  ferrocyanide 
Fe4[Fe(CN)6]3  has  a  brilliant  blue  color  ("Prussian  blue").  The 
ferrous  compound  Fe2Fe(CN)6,  or  perhaps  K2FeFe(CN)6,  is  white  but 
quickly  becomes  blue  by  oxidation.  The  ferrocyanides  are  not  poison- 
ous. 

Ferricyanide  of  potassium.  This  salt  is  easily  made  from  the 
ferrocyanide  by  oxidation : 

2K4Fe(CN)6  +  01,  ->  2KC1  +  2K8Fe(CN)6, 
or  2Fe(CN)6""  +  01,  -»  2Fe(CN)6'"  +  201'. 

It  forms  red  monosymmetric  prisms.  The  free  acid  H8Fe(CN)6  is  un- 
stable. Other  salts  may  be  prepared  by  precipitation.  Ferrous  ferri- 
cyanide  Fe3[Fe(CN)J2  is  deep-blue  in  color  («  TurnbulPs  blue  ").  With 
ferric  salts  only  a  brown  solution  is  obtained. 

Ferric  thiocyanate  Fe(CNS)3  is  formed  by  interaction  of  soluble 
thiocyanates  with  ferric  salts  (of.  p.  250).  It  is  deep-red  in  color  and 
gives  a  blood-red  solution  in  water.  Since  both  the  ions  are  colorless, 
the  solution  must  contain  chiefly  the  molecular  salt.  -  Its  formation 
furnishes  a  very  delicate  test  for  traces  of  ferric  salts. 

Iron  Carbonyls.  —  When  carbon  monoxide  is  led  over  finely 
divided  iron  at  40-80°,  or  under  eight  atmospheres  pressure  at  the 
ordinary  temperature,  volatile  compounds  of  the  composition  Fe(CO)4, 
the  tetracarbonyl,  and  Fe(CO)5,  the  pentacarbonyl,  are  formed.  When 


758  INORGANIC    CHEMISTRY 

the  gaseous  mixture  is  heated  more  strongly,  the  compounds  decom- 
pose again,  and  iron  is  deposited.  Illuminating-gas  burners  frequently 
receive  a  deposit  of  iron  from  this  cause. 

Ferrates.  —  A  red  solution  of  potassium  ferrate,  K2Fe04,  is  ob- 
tained by  passing  chlorine  through  caustic  potash  in  which  ferric 
hydroxide  is  suspended.  The  salt  crystallizes  in  red,  rhombic  prisms, 
isomorphous  with  the  sulphate  and  chromate  of  potassium.  It  alters 
quickly  in  solution,  in  consequence  of  hydrolysis  and  subsequent  de- 
composition of  the  ferric  acid,  depositing  ferric  hydroxide  and  giving 
off  oxygen.  Barium,  strontium,  and  calcium  salts  are  formed  as  red 
precipitates  by  double  decomposition. 

Analytical  Reactions  of  Compounds  of  Iron.  —  There  are  two 
ionic  forms  of  iron,  diferrion  Fe",  which  is  very  pale-green,  and  tri- 
ferrion  Fe***,  which  is  almost  colorless.  The  yellow  color  of  ferric 
salts  is  due  to  hydrolysis.  Ammonium  sulphide  forms,  with  both, 
black  ferrous  sulphide  which  is  soluble  in  dilute  acids.  The  hydrox- 
ides are  white  and  brown  respectively,  and  ferrous  carbonate  is  white. 
With  ferric  salts,  soluble  carbonates  yield  the  hydroxide.  With  ferro- 
cyanide  of  potassium,  ferrous  salts  give  a  white,  and  ferric  salts  a  blue, 
precipitate.  With  ferricyanide  of  potassium  the  former  give  a  deep- 
blue  precipitate,  and  the  latter  a  brown  solution.  Ferric  thiocyanate 
is  deep-red.  From  ferric  solutions  barium  carbonate  throws  down  fer- 
ric hydroxide.  When  sodium  acetate  is  added  in  excess  to  a  ferric 
salt,  a  red,  little  ionized,  but  much  hydrolyzed,  acetate  is  formed. 
When  the  solution  is  boiled  the  hydrolysis  is  increased,  and  an  insol- 
uble, basic  ferric  acetate  is  thrown  down.  With  borax,  iron  com- 
pounds give  a  bead  which  is  green  in  the  reducing  flame,  and  colorless 
or,  with  much  iron,  yellow  or  even  brown  when  oxidized. 

COBALT. 

The  Chemical  Relations  of  the  Element.  —  Cobalt  forms 
cobaltous  and  cobaltic  oxides  and  hydroxides  CoO  and  Co(OH)2,  Co208 
and  Co(OH)8,  respectively,  which  are  all  basic,  the  former  more  so 
than  the  latter.  The  cobaltous  salts  are  little  hydrolyzed,  but  the 
cobaltic  salts  are  completely  decomposed  by  water.  The  latter  also 
liberate  readily  one-third  of  the  negative  radical,  after  the  manner  of 
manganic  salts,  becoming  cobaltous.  Complex  cations  and  anions  con- 
taining cobalt  are  very  numerous  and  very  stable. 


IRON,    COBALT,    NICKEL  759 

Occurrence.  —  Cobalt  is  found  along  with  nickel  in  smaltite  CoAs2 
and  cobaltite  CoAsS.  The  pure  metal  may  be  made  by  Goldschmidt's 
process,  or  by  reducing  the  oxalate,  or  an  oxide,  with  hydrogen. 

Physical  and  Chemical  Properties.  —  The  metal  is  silver-white, 
with  a  faint  suggestion  of  pink.  It  is  less  tough  than  iron,  and  has 
no  commercial  applications.  It  displaces  hydrogen  slowly  from  dilute 
acids,  but  interacts  readily  with  nitric  acid. 

Cobaltous  Compounds.  —  The  chloride  may  be  made  by  treating 
the  oxide  with  hydrochloric  acid.  It  forms  red  prisms  of  CoCl2, 
6H20,  and  when  partially  or  completely  dehydrated  becomes  deep- 
blue.  Writing  made  with  a  diluted  solution  upon  paper  is  almost 
invisible,  but  becomes  blue  when  warmed  and  afterwards  takes  up 
moisture  from  the  air,  and  is  once  more  invisible  ("  sympathetic 
ink").  Most  cobaltous  compounds  are  red  when  hydrated  or  in  solu- 
tion (Co**)  and  blue  when  dehydrated.  The  blue  color  assumed  by  a 
strong  solution  of  cobaltous  chloride,  when  it  is  warmed,  or  when 
hydrochloric  acid  is  added  to  it,  is  explained  by  some  chemists  as 
being  due  to  repression  of  the  ionization  of  the  salt,  and  by  others 
as  being  due  to  the  formation  of  the  complex  anion  of  the  salt 
Co**.CoCl/'.  By  addition  of  sodium  hydroxide  to  a  cobaltous  salt,  a 
blue  basic  salt  is  precipitated.  When  the  mixture  is  boiled,  the  red 
hydroxide,  Co(OH)2  is  formed.  This  becomes  brown  through  oxidation 
by  the  air.  It  interacts  with  ammonium  hydroxide,  giving  a  soluble 
ammonio-cobaltous  hydroxide  (cf.  p.  661),  which  is  quickly  oxidized 
by  the  air  to  an  ammonio-cobaltic  compound  (see  below).  It  dissolves 
also  in  salts  of  ammonium  as  magnesium  hydroxide  does  (p.  644). 
When  dehydrated  it  leaves  the  black  cobaltous  oxide.  Cobaltous  sul- 
phate, CoS04,  7H20,  isomorphous  with  magnesium  sulphate,  and  the 
nitrate,  Co(N08)2, 6H20,  are  familiar  salts.  The  black  cobaltous  sul- 
phide, CoS,  is  precipitated  by  ammonium  sulphide  from  solutions  of 
all  salts,  and  even  by  hydrogen  sulphide  from  the  acetate,  or  a  solution 
containing  much  sodium  acetate  (cf.  p.  650).  Once  it  has  been  formed, 
it  does  not  interact  even  with  dilute  hydrochloric  acid,  having  appar- 
ently changed  into  a  less  active  form.  A  sort  of  cobalt  glass,  made 
by  fusing  sand,  cobalt  oxide,  and  potassium  nitrate,  forms,  when 
powdered,  a  blue  pigment  ("smalt")  used  in  china-painting  and  by 
artists. 


760  INORGANIC   CHEMISTRY 

Cobaltic  Compounds.  —  By  addition  of  a  hypochlorite  to  a  solu- 
tion of  a  cobaltous  salt,  cobaltic  hydroxide  Co(OH)3,  a  black  powder, 
is  precipitated.  Cautious  ignition  of  the  nitrate  gives  cobaltic  oxide, 
Co203.  Stronger  ignition  gives  the  commercial  oxide,  which  is  a  cobalto- 
cobaltic  oxide  Co8O4.  Cobaltic  oxide  dissolves  in  cold  hydrochloric 
acid,  but  the  solution  gives  off  chlorine  when  warmed.  By  placing 
cobaltous  sulphate  round  the  anode  of  an  electrolytic  cell,  crystals  of 
cobaltic  sulphate,  Co2(S04)8,  have  been  made  and  cobaltic  alums  have 
also  been  prepared. 

Complex  Compounds.  —  Potassium  cyanide  precipitates  from 
cobaltous  salts  a  brownish-white  cyanide  which  interacts  with  excess 
of  the  reagent,  giving  a  solution  of  potassium  cobaltocyanide  K4Co 
(CN)6  (cf.  p.  756).  This  compound  is  easily  oxidized  by  chlorine,  or 
even  when  the  solution  is  boiled  in  the  air,  and  the  colorless  potassium 
cobalticyanide  is  formed : 

4K4Co(CN)6  +  2H20  +  02  -»  4K8.Co(CN)6  +  4KOH. 

The  solution  gives  none  of  the  reactions  of  Co"*,  and  with  acids  the 
very  stable  cobalticyanic  acid,  H3Co(CN)3,  is  liberated. 

When  acetic  acid  and  potassium  nitrite  are  added  to  a  cobaltous  salt 
the  latter  is  oxidized  by  the  nitrous  acid  (liberated  by  the  acetic  acid) 
and  a  white  complex  salt  K8.Co(N02)6, 7iH20  (=  Co(N02)8,3KN02), 
potassium  cobaltinitrite,  is  thrown  down. 

Cobaltic  salts  give  with  ammonia  complex  compounds  which  are 
many  and  various.  The  cations  often  contain  negative  groups,  and 
are  such  as  Co(NHBy,  Co(NH3)6Cl"  and  Co(NH3)5N02".  Usually 
the  solutions  give  none  of  the  reactions  of  cobaltic  ions,  and  often  fail 
likewise  to  give  those  of  the  anion  of  the  original  salt. 

NICKEL. 

The  Chemical  Relations  of  the  Element. — Nickel  forms  nick- 
elous  and  nickelic  oxides  and  hydroxides  NiO  and  Ni(OH)2,  Ni208,  and 
Ni(OH)3,  but  only  the  former  are  basic.  The  nickelous  salts  resemble 
the  cobaltous  and  ferrous  salts,  but  are  not  oxidizable  into  correspond- 
ing nickelic  compounds.  Since  there  are  no  nickelic  salts,  there  are 
here  no  analogues  of  the  cobalticyanides  or  the  cobaltinitrites.  The 
complex  nickelous  salts,  like  the  complex  cobaltous  salts,  and  unlike 
the  complex  cobaltic  salts,  are  unstable,  and  so  give  some  of  the 
reactions  of  Ni**. 


IRON,    COBALT,    NICKEL  761 

\ 

Occurrence. — Nickel  occurs  free  in  meteorites  and  in  niccolite  NiAs 
and  nickel  glance  NiAsS.  The  chief  sources  at  present  are  garnierite 
(from  New  Caledonia)  and  pentlandite  [Ni,Fe]S  (from  Sudbury,  Ont.). 

Properties.  —  The  rnetal  is  white,  with  a  faint  tinge  of  yellow, 
is  very  hard,  and  takes  a  high  polish.  It  is  used  in  making  alloys, 
such  as  German  silver  (copper,  zinc,  nickel,  2:1:1)  and  the  "nickel" 
used  in  coinage  (copper,  nickel,  3:1).  Nickel  plating  on  iron  is 
accomplished  exactly  like  silver  plating  (p.  632).  The  bath  contains  an 
ammoniacal  solution  of  ammonium-nickel  sulphate  (NH4)2S04,  NiS04, 
6H20,  and  a  plate  of  nickel  forms  the  anode. 

The  metal  rusts  very  slowly  in  moist  air.  It  displaces  hydrogen 
with  difficulty  from  dilute  acids  but  interacts  with  nitric  acid. 

Compounds  of  Nickel.  —  The  chloride  NiCL,,  6H20,  is  made  by- 
treating  one  of  the  oxides  with  hydrochloric  acid,  and  is  green  in 
color  (when  anhydrous,  brown).  The  sulphate,  NiS04, 6H20,  which 
crystallizes  in  square  prismatic  forms  at  30-40°,  is  the  most  familiar 
salt.  The  heptahydrate  NiS04, 7H20,  obtained  from  cold  solutions,  is 
isomorphouswith  magnesium  sulphate.  Nickelous  hydroxide,  Ni(OH)2, 
is  formed  as  an  apple-green  precipitate,  and  when  heated  leaves  the 
green  nickelous  oxide,  NiO.  It  dissolves  in  ammonium  hydroxide, 
giving  a  complex  nickel-ammonia  cation.  It  is  soluble  also  in  salts  of 
ammonium  (cf.  p.  661).  By  cautious  ignition  of  the  nitrate,  nickelic 
oxide,  NijOg,  is  formed  as  a  black  powder.  The  oxides  and  salts,  when 
heated  strongly  in  oxygen  give  the  oxide  Ni3O4.  The  last  two  oxides 
liberate  chlorine  when  treated  with  hydrochloric  acid,  and  give  nickel- 
ous chloride.  Nickelic  hydroxide,  Ni(OH)3,  is  a  black  precipitate 
formed  when  a  hypochlorite  is  added  to  any  salt  of  nickel.  Nickelous 
sulphide  is  thrown  down  by  ammonium  sulphide,  and  behaves  like 
cobaltous  sulphide  (p.  759).  It  forms  a  brown  colloidal  solution  when 
excess  of  the  precipitant  is  used,  and  is  then  deposited  very  slowly. 

With  potassium  cyanide  and  a  salt  of  nickel  the  greenish  nickelous 
cyanide,  Ni(CN)2,  is  first  precipitated.  This  dissolves  in  excess  of  the 
reagent,  and  a  complex  salt  K2Ni(CN)4,  H2O(  =  2KCN,Ni(CN)2)  may 
be  obtained  from  the  solution.  This  salt  is  of  different  composi- 
tion from  the  corresponding  compounds  of  cobalt  and  iron,  and  is  less 
stable.  Thus,  with  bleaching  powder,  it  gives  Ni(OH)3  as  a  black 
precipitate.  When  the  solution  is  boiled  in  the  air  no  oxidation  to  a 
complex  nickelicyanide  occurs,  and  indeed  no  such  salts  is  known. 


762  INORGANIC   CHEMISTRY 

This  fact  enables  the  chemist  to  separate  cobalt  and  nickel,  for  when 
the  mixed  cyanides  are  boiled  and  then  treated  with  bleaching  powder, 
the  cobalticyanide  is  unaffected.  With  potassium  nitrite  and  acetic 
acid  no  insoluble  compound  corresponding  to  that  given  by  cobalt  salts 
is  formed  by  salts  of  nickel.  The  only  known  compound  which  could 
be  formed,  4KN02,Ni(]Sr02)2,  is  soluble.  This  action  also  is  used  for 
the  purpose  of  separation. 

When  finely  divided  nickel,  made  by  reducing  the  oxide  or  oxalate 
with  hydrogen  at  a  moderate  temperature,  is  exposed  to  a  stream  of 
cold  carbon  monoxide,  nickel  carbonyl  Ni(CO)4  is  formed.  This  is  a 
vapor  and  is  condensable  to  a  colorless  liquid  (b.-p.  43°  and  m.-p.  —25°). 
The  vapor  is  poisonous.  When  heated  to  150-180°  it  is  dissociated  and 
nickel  is  deposited.  Cobalt  forms  no  corresponding  compound. 

Analytical  Reactions  of  Compounds  of  Cobalt  and  Nickel.  — 

The  cobalt  ion  Co**  is  pink,  and  the  nickelous  ion  Ni**  green.  The 
reactions  used  in  analysis  have  been  described  in  the  preceding  para- 
graphs. With  borax,  cobalt  compounds  give  a  blue  bead,  and  nickel 
compounds  a  bead  which  is  brown  in  the  oxidizing  flame  and  cloudy, 
from  the  presence  of  gray,  metallic  nickel,  when  reduced. 

Exercises.  —  1.  What  would  be  the  interactions  of  calcium  car- 
bonate when  fused  with  sand  and  with  clay,  respectively  ? 

2.  Make   equations   representing    (a)    the    oxidation    of    ferrous 
chloride  by  air,  (b)  the  hydrolysis  of  ferrous  carbonate  and  the  oxida- 
tion of  ferrous  hydroxide,  (c)  the  oxidation  of  ferrous  sulphate  with 
excess  of  sulphuric  acid  by  hypochlorous  acid,  (d)  the  formation  of 
ferrous  and  ferric  tannates  (p.  754),  (e)the  reduction  of  ferric  chloride 
by  iron,  by  hydrogen  sulphide,  and  by  stannous  chloride,  respectively, 
(f)  the  dry  distillation  of  basic  ferric  sulphate,  (g)  the  formation  of 
ferric  ferrocyanide  and  of  ferrous  ferricyanide,  (h)  the  hydrolysis  and 
decomposition  of  potassium  ferrate. 

3.  Explain  the  solubility  of  cobaltous  and  nickelous  hydroxide  in 
salts  of  ammonium. 

4.  Construct  equations  to  show  the  formation  (a)  of  the  insoluble 
potassium   cobaltinitrite  (nitric  oxide   is   given    off),   (b)  of   nickelic 
hydroxide  from  nickelous  chloride  and  sodium  hypochlorite. 

5.  Tabulate  in  detail  the  chemical  relations  of  the  elements  cobalt 
and  nickel,  with  especial  reference  to  showing  the  resemblances  and 
differences. 


CHAPTER  XLV 

THE   PLATINUM  METALS 

THE  remaining  elements  of  Mendele Jeff's  eighth  group  divide  them- 
selves into  two  sets  of  three  each.  Just  as  iron,  cobalt,  and  nickel 
have  similar  atomic  weights  and  much  the  same  specific  gravity, 
(7.8-8.8),  so  ruthenium  (Ru,  at.  wt.  101.7),  rhodium  (Rh,  at.  wt.  103), 
and  palladium  (Pd,  at.  wt.  106.5)  have  specific  gravities  from  12.26  to 
11.5.  Similarly  osmium  (Os,  at.  wt.  191),  iridium  (Ir,  at.  wt.  193), 
and  platinum  (Pt,  at.  wt.  194.8)  form  a  triad  with  specific  gravities 
from  22.5  to  21.5.  Chemically,  ruthenium  shows  the  closest  resem- 
blance to  osmium,  and  both  are  allied  to  iron.  Similarly,  rhodium  and 
iridium,  and  palladium  and  platinum  are  natural  pairs. 

The  six  elements  are  found  alloyed  in  nuggets  and  particles  which 
are  separated  from  alluvial  sand  by  washing.  Platinum  forms  60-84 
per  cent  of  the  whole.  The  chief  deposits  are  in  the  Ural  Mountains, 
smaller  amounts  being  found  in  California,  Australia,  BorneOj  and 
elsewhere.  The  components  are  separated  by  a  complex  series  of 
chemical  operations. 

Ruthenium  and  Osmium.  —  These  metals  are  gray  like  iron, 
while  the  other  four  are  whiter  and  more  like  cobalt  and  nickel.  They 
also  resemble  iron  in  being  the  most  infusible  members  of  their 
respective  sets.  Both  melt  considerably  above  2000°.  They  likewise 
resemble  iron  in  uniting  easily  with  free  oxygen,  while  the  other  four 
elements  do  not.  Ruthenium  gives  RuO2  and  even  Ru04,  although 
the  latter  oxide  is  more  easily  obtained  indirectly.  Osmium  gives 
Os04, "  osmic  acid,"  a  white  crystalline  body  melting  at  40°  and  boiling 
at  about  100°.  The  odor  and  irritating  effects  of  the  vapor  recall 
chlorine  (G-k.,  007x17,  odor).  The  substance  is  not  an  acid,  or  even  an  acid 
anhydride.  The  aqueous  solution  is  used  in  histology,  and  stains 
tissues  in  consequence  of  its  reduction  by  organic  bodies  to  metallic 
osmium.  It  is  affected  particularly  by  fat.  Osmic  acid  also  hardens 
the  material  without  distorting  it.  It  will  be  observed  that  ruthenium 
and  osmium  have  a  maximum  valence  of  eight. 

These  elements  resemble  iron  in  giving  ruthenates  and  osmates, 

763 


764  INORGANIC   CHEMISTRY 

like  K2Ru04  and  K20s04,  but  no  corresponding  oxide  or  acid.  Potas- 
sium rutheiiate  resembles  potassium  manganate  and  gives,  when 
diluted,  an  oxide  of  ruthenium  and  potassium  perruthenate  KRu04. 
There  are  also  salts  corresponding  to  the  oxides  Ru208  and  Os203. 

Rhodium  and  Iridium.  —  These  metals  are  not  attacked  by 
aqua  regia,  while  the  other  four  are  dissolved,  more  or  less  slowly. 
They  are  harder  than  platinum,  and  iridium  is  alloyed  with  this  metal 
for  the  purpose  of  increasing  its  resistance  to  the  action  of  acids. 
They  resemble  cobalt  in  having  no  acid-forming  properties.  Salts 
corresponding  to  the  oxides  Kh203  and  Ir208  are  formed,  and  iridium 
gives  compounds  derived  from  Ir02  as  well.  The  most  familiar  com- 
pounds of  iridium  are  the  complex  chlorides  X8IrCl6  (=  3XCl,IrCl8) 
and  XjjIrClg  (=  2XCl,IrCl4).  The  solutions  of  the  latter  are  red,  and 
the  acid,  chloroiridic  acid  H2IrCl6,  is  often  found  in  commercial  chloro- 
platinic  acid  H2PtCl6,  and  confers  upon  it  a  deeper  color. 

Palladium  and  Platinum.  —  Palladium  is  the  only  metal  of 
this  family  which  is  attacked  by  nitric  acid.  Palladium  and  platinum 
form  -ous  and  -ic  compounds  of  the  forms  PdX2  and  PdX4  respectively. 
The  oxides  PdO  and  PtO  and  corresponding  hydroxides  are  basic. 
When  quadrivalent,  the  metals  appear  chiefly  in  complex  compounds, 
like  H2.Pt016,  Hjj-PdClg,  in  which  the  metal  is  in  the  anion.  Platinum 
gives  also  platinates  derived  from  the  oxide  PtO2,  and  quadrivalent 
platinum  furnishes  no  well  defined  salts  in  which  it  constitutes  by 
itself  the  positive  ion. 

Palladium.  —  This  metal,  named  from  the  planetoid  Pallas,  is 
noted  chiefly  for  its  great  tendency  to  absorb  hydrogen.  At  1000°  it 
takes  up  about  650  times  its  own  volume.  The  amount  absorbed 
varies  continuously  with  the  concentration  (pressure)  of  the  hydrogen, 
although  not  according  to  a  uniform  rule,  and  the  case  is  therefore  re- 
garded as  being  one  of  solid  solution.  When  a  strip  of  palladium  is 
made  the  cathode  of  an  electrolytic  cell,  over  900  volumes  of  hydrogen 
may  be  occluded.  This  absorbed  hydrogen,  in  consequence  of  the  cat- 
alytic influence  of  the  metal,  reacts  more  rapidly  than  does  the  gas,  and 
consequently  a  strip  of  hydrogenized  palladium  will  quickly  precipitate 
copper  and  other  metals  less  electro-positive  than  hydrogen  and  will 
reduce  ferric  and  other  reducible  salts  : 

CuS04  +  H2  ->  H2S04  +  Cu,   or   Cu"+  H2  -»  2H'  -f  Cu. 
2FeC^  +  H,  -»  2FeCLj  +  2HC1,   or   2Fe'"  +  H2->2Fe"  +  2H. 


THE   PLATINUM   METALS  765 

The  palladious  salts  are  such  as  PdCLj,  PdS04,  Pd(N08)2.  Palladic 
chloride,  formed  by  treating  the  metal  with  aqua  regia,  is  contained  in 
the  solution  in  the  form  of  chloropalladic  acid  HgPdClg  (=  2HC1, 
PdCl4)  and  gives  difficultly  soluble  salts  like  K^dC^.  When  the 
solution  of  the  acid  is  boiled,  however,  chlorine  is  given  oft'  and  PdCl2 
or  H2PdCl4  remains  in  solution. 

Platinum.  —  This  metal  (dim.  of  Sp.plata,  silver)  is  grayish-white 
in  color,  and  is  very  ductile.  At  a  red  heat  it  can  be  welded.  It  does 
not  melt  in  the  Bunsen  flame  buf  fuses  easily  in  the  oxy hydrogen  jet. 
On  account  of  its  very  small  chemical  activity  it  is  used  in  electrical 
apparatus  and  for  making  wire,  foil,  and  crucibles  and  other  vessels 
for  use  in  laboratories.  It  unites,  however,  with  carbon,  phosphorus, 
and  silicon,  becoming  brittle,  and  forms  fusible  alloys  with  metals  like 
antimony  and  lead.  Hence  care  has  to  be  taken  not  to  heat  in  vessels 
made  of  it  compounds  from  which  these  elements  may  be  liberated.  It 
also  interacts  with  fused  alkalies,  giving  platinates,  but  the  alkali  car- 
bonates may  be  melted  in  vessels  of  platinum  with  impunity.  The 
oxygen  acids  are  without  action  upon  it,  but  the  free  chlorine  in  aqua 
reyia  converts  it  into  chloroplatinic  acid  H2PtCl6. 

The  metal  condenses  oxygen  upon  its  surface  and  it  dissolves 
hydrogen.  The  finely  divided  forms  of  the  metal,  such  as  platinum 
•ponge  made  by  igniting  ammonium  chloroplatinate  (NH4)2PtCl6,  and 
platinum  black  made  by  adding  zinc  to  chloroplatinic  acid,  show  this 
behavior  very  conspicuously.  They  cause  instant  explosion  of  a  mix- 
ture of  oxygen  and  hydrogen,  in  consequenc  of  the  heat  developed  by 
the  rapid  union  of  that  part  of  the  gases  which  is  condensed  in  the 
metal.  A  heated  spiral  of  fine  platinum  wire  will  continue  to  glow  if 
immersed  in  the  mixture  of  alcohol  vapor  and  oxygen  formed  by  lead- 
ing oxygen  through  liquid  alcohol.  The  heat  is  developed  by  the  inter- 
action which  takes  place  between  the  substances  with  great  speed  at  the 
surface  of  the  platinum.  Platinum  sponge  is  the  active  constituent  of 
the  contact-mass  used  in  making  sulphur  trioxide  (p.  380). 

Platinum  is  the  only  otherwise  suitable  substance  which  has  the 
same  coefficient  of  expansion  as  glass,  and  it  is  consequently  fused  into 
incandescent  bulbs  and  furnishes  the  electrical  connection  with  the 
filament  in  the  interior.  Large  amounts  are  also  consumed  in  photo- 
graphy. The  price  of  the  metal  is  subject  to  great  variations,  since  a 
rainy  season  in  the  Caucasus  will  render  larger  amounts  accessible  to 
the  miners,  but,  on  the  whole,  the  many  applications  which  have  been 
found  for  it  have  tripled  its  price  in  the  last  twenty  years. 


766  INORGANIC   CHEMISTRY 

When  special  resistance  to  chemical  or  mechanical  influences  is 
required,  as  in  standard  meters  for  international  reference,  or  points 
of  fountain  pens,  the  alloy  with  indium  is  employed  (cf.  p.  240). 

Compounds  of  Platinum.  —  Flatinous  chloride  is  made  by  pass- 
ing chlorine  over  finely  divided  platinum  at  240-250°  or  by  heating 
chloroplatinic  acid  to  the  same  temperature.  It  is  greenish  and  in- 
soluble in  water,  but  forms  with  hydrochloric  acid  the  soluble  chloro- 
platinous  acid  H2PtCl4.  The  potassium  salt  of  this  acid  K2PtCl4  is 
used  in  making  platinum  prints  (cf.  p.  635).  Bases  precipitate  black 
platinous  hydroxide  Pt(OH)2  which  interacts  with  acids  but  not  with 
bases.  Gentle  heating  gives  the  oxide  PtO  and  stronger  heating  the 
metal.  With  potassium  cyanide  and  barium  cyanide  soluble  platino- 
cyanides,  K2Pt(CN)4, 3H20  and  BaPt(CN)4, 4H20,  are  formed.  These 
substances,  when  solid,  show  strong  fluorescence,  converting  X-rays  as 
well  as  ultra-violet  rays  into  visible  radiations.  The  barium  salt  is 
used  to  coat  screens  on  which  the  shadows  cast  by  X-rays  are  received. 

Platinic  chloride  PtCl4  may  be  made  by  treating  the  metal  with 
aqua  regia  and  heating  the  chloroplatinic  acid  H2PtCl6  so  formed  in 
a  stream  of  chlorine  at  360°.  When  dissolved  in  water,  however,  it 
gives  a  compound  H2PtCl40  in  which  platinum  is  in  the  anion.  Its 
solution  deposits  red,  non-deliquescent  crystals  of  H2PtCl40, 4H20. 

Chloroplatinic  acid  forms  reddish-brown  deliquescent  crystals 
H2PtCl6,6H20.  With  potassium  and  ammonium  salts,  it  yields  the 
sparingly  soluble,  yellow  chloroplatinates  KgPtClg  and  (NH4)2PtCl6 
(cf.  p.  561),  in  solutions  of  which  the  platinum  migrates  towards  the 
anode  and  silver  salts  precipitate  Ag2PtCl6  and  not  silver  chloride. 

Bases  interact  with  chloroplatinic  acid,  giving  a  yellow  or  brown 
precipitate  of  platinic  hydroxide  Pt(OH)4.  This  substance  interacts 
both  with  acids  and  with  bases.  In  the  former  case  the  salts,  which 
probably  are  formed,  have  not  been  isolated.  In  the  latter  case  plati- 
natea,  like  Na^H^PtgO^,  H20,  have  been  obtained.  Both  sets  of  plati- 
num compounds  interact  with  hydrogen  sulphide,  giving  the  sulphides, 
PtS  and  PtS2  respectively.  These  are  black  powders  which  dissolve 
in  yellow  ammonium  sulphide  solution,  much  as  do  the  sulphides  of 
gold,  arsenic,  and  other  metals,  giving  ammonium  sulphoplatinates. 

There  are  numerous  complex  compounds  with  ammonia  represent- 
ing both  the  platinous  and  platinic  series.  They  resemble  the 
ammonio-cobaltic  compounds  in  their  behavior. 


INDEX 


V  Acids  are  all  listed  under  "  acid  "  and  salts  under  the  positive  radical. 


ACCUMULATORS,  702 
Acetylene,  496,  515 
Acid,  acetic,  498 

ionization  of,  578 

antimonic,  716 

arsenic,  711 

arsenious,  711 

boracic,  527 

boric,  527 

bromic,  277 

carbolic,  441 

carbonic,  481 

chloric,  274 

chlorosulphuric,  399 

chlorous,  275 

chromic,  728 

disulphuric,  388 

dithionic,  398 

formic,  497 

hydrazoic,  422,  452 

hydriodic,  239,  see  Hydrogen  iodide 

hydrobromic,     233,     see     Hydrogen 
bromide 

hydrochloric,  93,  see  Hydrogen  chlo- 
ride 

chemical  properties,  185 
of  constant  boiling-point,  182 

hydrocyanic,  507 

hydrofluoboric,  527 

hydrofluoric,  242 

hydrofluosilicic,  521 

hypobromous,  276 

hypochlorous,  176,  267 

hyponitrous,  450 

hypophosphorous,  469 

hyposulphurous,  393 

iodic,  277 

metaphosphoric,  465,  468 

metastannic,  696 

muriatic,  see  Acid,  hydrochloric 


Acid,  nitric,  93,  438 

fuming,  440 

oxidizing  actions  of,  446 

preparation,  439 

properties,  439 
nitrosylsulphuric,  383 
nitrous,  449 

orthophosphoric,  465,  466,  732 
oxalic,  499 
palmitic,  505 
pentathionic,  398 
perchloric,  276 
perchromic,  729 
periodic,  278 
permanganic,  743 
persulphuric,  397 
phosphorous,  469 
picric,  441 
pyrophosphoric,  467 
tannic,  754 
selenic,  402 
silicic,  523 
a-stannic,  696 
sulphuric,  93,  382 

chamber  process,  383 

constitution,  391 

contact  process,  380 

fuming,  388 

Nordhausen,  388 

properties,  387,  389 

uses,  393 
sulphurous,  393 
telluric,  403 
tetraboric,  528 
tetrathionic,  398 
thiosulphuric,  396 
trithionic,  398 
Acidimetry,  351 
Acids,  70,  92,  119,  281,  345 

activity  of,  see  Chemical  activity 


767 


768 


INDEX 


Acids,  complex,  salts  of,  360,  363,  536 

constitution  of,  470 

dibasic,  ionization  of,  346,  374 

ionic  double  decompositions  of,  349 

polythionic,  398 
Acker  process,  169,  570 
Adsorption,  476 
Affinity,  constant,  254 

criticism  of  word,  29 

inapt  use  of  word,  110 

see  Chemical  activity 
Agar-agar,  313 
Air,  a  mixture,  430 

composition,  431 

liquid,  434 

solubility,  155 
Albumins,  415 
Alcohol,  ethyl,  503 

methyl,  503 
Alizarine,  689 
Alkalies,  348,  548 
Allotropic  modifications,  459 
Alloys,  532 
Alumina,  686 
Aluminates,  685 
Aluminium,  682 

acetate,  688 

carbide,  494 

chloride,  684 

hydroxide,  684 

oxide,  686 

sulphate,  686 

sulphide,  687 
Aluminothermy,  540 
Alums,  547,  632,  687,  730,  741,  756 
Amalgams,  532 
Amethyst,  522 
Ammonia,  417 
Ammonia-soda  process,  574 
Ammonion,  568 
Ammonium  amalgam,  567 

bromide,  564 

carbamate,  566 

carbonate,  565 

chloride,  564 

compounds,  420,  564 

cyanate,  488,  566 

hydroxide,  565 

iodide,  564 

nitrate,  451,  565 


Ammonium  sesquicarbonate,  566 

sulphate,  566 

sulphides,  566 

sulphostannate,  697 

thiocyanate,  566 
Amorphous  substances,  123,  139,  369, 

-       475,  519 

Analysis,  application  of  ionic  hypothe- 
sis in,  343 

of  compounds  of  carbon,  623 

recognition    of    cations    ("  metals") 

in,  660 
Analytical  reactions,  aluminium,  691 

ammonium,  568 

antimony,  720 

arsenic,  720 

barium,  612 

bismuth,  720 

cadmium,  652 

calcium,  608,  612 

chromium,  731 

cobalt,  762 

copper,  625 

gold,  639 

iron,  758 

lead,  705 

manganese,  745 

magnesium,  645 

mercury,  659 

nickel,  762 

potassium,  561 

silver,  635 

sodium,  577 

strontium,  609,  612 

tin,  697 

zinc,  650 
Anhydride,  chromic,  728 

hypochlorous,  267 

iodic,  278,  279 

nitric,  440 

nitrous,  450 

perchloric,  276 

permanganic,  743 

phosphoric,  464 

sulphuric,  380 

sulphurous,  378 
Anhydrides,  71 

mixed,  397,  463 

various  acids  from  one,  278 
Anions,  322 


INDEX 


769 


Anode,  322 
Antichlor,  397 
Antimonites,  716 
Antimony,  713 

nitrate,  716 

oxychloride,  715 

pentachloride,  715 

pentoxide,  715 

sulphate,  716 

sulphides,  716 

trichloride,  715 

trioxide,  715 

Antimonyl  compounds,  716 
Apatite,  455 

Aqueous  tension,  82,  88,  116,  135 
Aqua  regia,  448 
Aragonite,  590 

Argentic  compounds,  see  Silver 
Argol,  501 
Argon,  428,  434 
Arsenic,  707 

pentoxide,  711 

sulphides,  712 

trichloride,  710  ' 

trioxide,  710 
Arsine,  709 

decomposition  of,  252 
Asbestos,  642 
Asphalt,  492 
Assaying,  639 
Association,  242,  294 
Atmosphere,  426 
Atom,  193 

definition  of,  218 
Atomic  hypothesis,  217 
Atomic  volume,  407 
Atomic  weight,  definition  of,  200 
Atomic  weights,  50 

determination  of,  196 

table  of,  inside  rear  board 
Attributes  of  bodies,  35 
Avogadro's  hypothesis,  131,  190,  288 
Azurite,  616 

BACTERIA,  nitrifying,  438 
Balanced  actions,  248 
Barium,  609 

carbonate,  610 

chlorate,  611 

chloride,  611 


Barium,  chloride,  hydrated,  121 

chromate,  728 

dioxide,  611,  see  Peroxide 

hydroxide,  612 

nitrate,  612 

oxide,  611 

peroxide,  63,  303,  304,  611 
dissociation  of,  257 

sulphate,  610 

sulphide,  610 
Baryta-water,  612 
Bases,  71,  119,  281,  348 

ionic  double  decompositions  of,  350 

solubilities  of,  544 
Basicity,  358 

Battery-cells,  666,  see  Cells 
Battery,  theory  of,  667 
Bauxite,  685 

Bead  tests,  468,  528,  731,  745 
Beer,  502 
Benzene,  497 
Benzine,  492 

Beryllium,  see  Glucinum,  641 
Bessemer  process,  749 
Bicarbonates,  see  Hydrogen  carbonates 
Bichromates,  see  Bichromates 
Bismuth,  717 

compounds,  718 

Bisulphates,  see  Hydrogen  sulphates 
Bisulphites,  see  Hydrogen  sulphites 
Bleaching,  269,  394 

powder,  266,  602 

Blue-stone  (blue  vitriol),  120,  625 
Body,  meaning  of  term,  34 
Boiling-point  of  mixed  liquids,  182,  503 
Boiling-points,  elevation  of,  162 

of  solutions,  162 
Bone-ash,  456,  605 
Bone  black,  476 
Borates,  528 
Borax,  528 
Boron,  526 

carbide,  529 

nitride,  529 

trichloride,  527 

trifluoride,  527 

trioxide,  529 
Boyle's  law,  84,  288 
Brandy,  502 
Brass,  619 


770 


INDEX 


Brimstone,  367 

Erin's  oxygen  process,  63,  611 

Bromine,  227 

chemical  properties,  230 

physical  properties,  229 

preparation,  227 
Bronze,  619 
Brucite,  644 

CADMIUM,  650 

compounds,  651 
Caesium,  563 
Calamine,  646 
Calcination,  456 
Calcion,  608 
Calcite,  590 
Calcium,  588 

carbide,  478,  602 

carbonate,  590 

chloride,  589 

chromate,  727 

fluoride,  590 

hydride,  588 

hydrogen  sulphite,  395 

hydroxide,  595 

nitrate,  602 

oxalate,  596 

oxide,  594 

phosphates,  605 

phosphide,  461 

silicate,  606 

sulphate,  603 

sulphide,  604 

Calculations,  58,  65,  91,  210 
Calomel,  654 
Calorie,  76 
Cane-sugar,  500 
Carbohydrates,  500 
Carbon,  473 

amorphous,  475 

properties,  477 

dioxide,  479 
properties,  480 

disulphide,  488 

monoxide,  485 

properties,  486 
Carbonates,  482 

Carbonates,    acid,   see   Hydrogen   car- 
bonates 
Carbonyl  chloride,  487 


Carborundum,  520 

Carnallite,  551,  642,  643 

Cassiterite,  693 

Catalytic  action,  75,  96,  111,  5«8 

Cathode,  322 

Cations,  322 

Cause,  in  Science,  30 

Caustic  potash,  552 

Cell,  Bunsen,  673 

Daniell,  672 

dichromate,  673 

dry,  673 

gravity,  672 

Leclanche",  673 

storage,  702 
Cells,  battery,  666 

concentration,  673 
Cellulose,  500 
Cement,  596 
Cerium,  510,  682,  705 
Chalk,  590 

Chalybeate  water,  483 
Chance's  process,  574 
Characteristics    of    chemical    phenom- 
ena, 5,  16,  17,  26,  27,  41,  48 
Charcoal,  476 

properties,  476 
Charles'  law,  85,  87 
Chemical  activity,  cause  of,  28 

measurement  of,  28,   79,   111,  678, 
679 

of  acids,  347,  356,  679 
Chemical  change,   spontaneously  pro- 
gressing, 27 

varieties  of,  14,  99,  121,  176, 181,  187, 

360  (ionic),  397,  520 
Chemical    relations   of   elements,    177, 

226,  719 

Chemistry,  methods  of  work  in,  36 
Chinese  white,  648 
Cinnabar,  653 
Chlorates,  272 
Chlorides,  electrolysis  of,  169 

interaction  with  acids,  179 

solubility,  185 
Chlorine,  168 

bleaching  action,  177 
chemical  properties,  174 
chemical  relations,  177 

dioxide,  275 


INDEX 


771 


Chlorine,  physical  properties,  173 

preparation,  168 

test  for,  175 

uses,  177 
Chloroform,  494 
Choke-damp,  493 
Chrome-alum,  730 
Chromic  acetate,  731 

anhydride,  728 

chloride,  729 

oxide,  730 

hydroxide,  729 

sulphate,  730 
Chromites,  729 
Chromium,  722 

trioxide,  728 

Chromous  compounds,  731 
Chromyl  chloride,  728 
Clay,  525,  689 
Coal,  477 
Cobalt,  758 

complex  compounds,  760 

zincate,  648 

Cobaltic  compounds,  760 
Cobaltous  compounds,  759 
Coke,  475 

Colloidal  solution,  523 
Colloids,  524 
Columbium,  721 
Combination,  14 
Combining  proportions,    measurement 

of,  43 
Combining  weights,  46 

law  of,  48 
Combustion,  71 

Common  ion,  effect  of  adding,  580 
Complex  cations,  534,  614,  622,  661,  662 

cadmium,  652 

chromium,  729,  730 

cobalt,  759,  760 

copper,  620,  622,  623,  625 

mercury,  659 

nickel,  761 

platinum,  766 

silver,  629,  631 

zinc,  648 

Component,  definition  of,  34 
Compound,  15 
Compounds,  binary,  103 

molecular,  123,  443 


Concentration,  molar,  250 

molecular,  law  of,  252 
illustration  of  law  of,  394 

of  gases,  80 
Conditions,  35,  52 
Conductivity,  electrical,  325- 
Condy's  disinfecting  fluid,  743 
Congo  red,  355,  689 
Conservation  of  mass,  17 
Constant,  depression,  291 

ionization,  297 

heat  summation,  law  of,  78 
Constituent,  definition  of,  34 
Constitution  of  acids,  470 

of  substances,  104,  224,  279,  308,  391, 

421,  441,  451,  466 
Copper,  615 

acetylene,  496,  516 

properties,  618 
Corpuscles,  222,  735 
Couples,  electrical,  96,  673 
Corrosive  sublimate,  655 
Crayon,  590 
Cream  of  tartar,  502 
Critical  phenomena,  133 
Cryolite,  683 
Crystal,  forms,  137 

structure,  140 

Crystallization,  water  of,  123,  181 
Cupellation,  627,  639,  700 
Cupric  acetate,  624 

bromide,  621 

ionization  of,  335 

carbonate  (basic),  623 

chloride,  619 

cyanide,  624 

ferrocyanide,  286,  299,  757 

hydroxide,  623 

iodide,  621 

nitrate,  623 

oxide,  43,  622 

sulphate,  624 

hydrated,  120,  122 

sulphide,  625 
Cuprous  chloride,  620 

cyanide,  624 

iodide,  621 

oxide,  622 

sulphide,  625 
Cyanogen,  507 


772 


INDEX 


D  ALTON'S  law  of  partial  pressures,  88 
Deacon's  process,  170 
Decomposition,  15 

potentials,  324,  675 
Decrepitation,  556 
Definite  proportions,  law  of,  41 
Deliquescence,  162 
Densities  of  gases,  192 

measurement  of,  89,  90 
Densities  of  solutions,  164 
"Derived  from,"  meaning  of  terra,  278 
Dextrose,  500 
Dialysis,  523 
Diamond,  474 
Dichromates,   chemical   properties   of, 

725 
Diffusion,  in  gases,  107 

in  liquids,  136 

Dilution  formula,  Ostwald's,  298 
Dimorphous  substances,  369 
Discharging  potentials,  324,  675 
Displacement,  95,  99 

ionic,  342,  361 
Dissociation,  121 

hydrolytic,  181,  see  Hydrolysis 

in  solution,  proofs  of,  281,  289,  292, 
293 

of  compounds  (by  heat),  213 

pressure,  257 
Distillation,  38 

fractional,  399,  492 

in  vacuo,  276,  471 
Dolomite,  642 
Double  decomposition,  15 
Dulong  and  Petit 's  law,  201 
Dyeing,  688 

EARTHENWARE,  689 
Efflorescence,  121,  123,  162 
Electrolysis,  310,  675 

of  chlorides,  169 
Electrolytes,  296,  310 
Electromotive  chemistry,  664 

force,  665 

series,  361,  670,  676 
Electrons,  222,  735 
Electroplating,  632 
Element,  definition  of,  31,  32 

not  same  as  simple  substance,   32, 
57 


Elements,  base-forming,  see  Elements, 
metallic 

chemical  relations  of,  177,  226,  719 

metallic,  119,  337,  404,  533 

classification  of,  537 
Elements,  molecular  weights  of,  204 

negative,  322,  405 

non-metallic,  119,  337,- 405,  534 

positive,  322,  404 

quantities  of  terrestrial  material,  33 
Endothermal  actions,  27 
Energy,  and  chemical  change,  19,  28 

chemical,  25 

conservation  of,  23 

definitions  of,  23,  24 

factors  of,  677 

free  or  available,  26 

units  employed  in  measuring,  25 
Enzymes,  501 
Epsom  salts,  644 
Equations,  54 

making  of,  55,  69,  209,  274,  447 

molecular,  208 

Equilibria,  displacement  of,  257 
Equilibrium,  117,  135 

chemical,  246 

ionic,  297 

considered  quantitatively,  578 
Equivalents,  49,  102 
Esters,  504 
Ether,  147,  506 

luminiferous,  24 
Ethyl  acetate,  504 

sulphate,  504 
Ethylene,  495 

bromide,  496 

Excess,  meaning  of  term,  43 
Exothermal  actions,  27 
Explanation,  its  nature,  10 
Explosives,  452 

FACTORS  of  energy,  77,  677 

Fat,  505 

Fehling's  solution,  623 

Fermentation,  501 

Ferrates,  758 

Ferric,  compounds,  754,  755,  756 

thiocyanate,  250,  757 
Ferrous,  compounds,  753,  754 

sulphide,  12,  42,  371 


INDEX 


773 


Fertilizers,  415,  467,  605 
Filtrate,  11 
Fire-damp,  493 
Flame,  509 

Bunsen,  511,  512 
Flame,  chemical  changes  in,  513 
Flash-light  powder,  643 
Fluorine,  preparation,  240 

properties,  241 
Fluorite,  590 
Flux,  529 
Formulae,  53 

graphic,  see  Constitution  of  substances 

making  of,  55,  see  Equations 

molecular,  of  compounds,  203 
of  simple  substances,  204 

structural,  see  Constitution  of  sub- 
stances 

Franklinite,  646,  756 
Freezing  mixtures,  164 
Freezing-point,     depression     constant, 

291,  294 
Freezing-points,  depression  of,  163,  290 

of  solutions,  163,  290 
Furnace,  blast,  748 

electric,  457,  479 

reverberatory,  572 

GALENA,  698,  705 

Gallium,  681 

Garnet,  525,  683 

Gas,  illuminating,  composition  of,  513 

Gases,  densities  of,  192 

drying  of,  100 

liquefaction  of,  432 

purification  of,  100 

solubility  of,  153 
Gasoline,  492 
Gay-Lussac's  law,  see  Charles'  law 

of  combining  volumes,  125 
Germanium,  692 
Glass,  606 

Glauber's  salt,  120,  158,  160,  576 
Glucinum,  413,  641 
Glucose,  500 
Glycerine,  503 
Gold,  635,  638 

Gram-molecular  volume,  195 
Grape-sugar,  500 
Graphite,  475 


Green  fire,  611 
Greenockite,  650 
Gun-cotton,  441,  452,  504 
Gunpowder,  38,  557 
Gypsum,  603 

HALOGEN  family,  226,  243,  279 
Heat,  of  fusion,  115 
of  solution,  164 
of  vaporization,  118 
Heavy-spar,  610 
Helium,  436,  563,  736 
Homologous  series,  491 
Hydrargyllite,  685 
Hydrates  (hydroxides),  120 

(substances     containing     water     of 

crystallization),  120 
of  salts,  composition  of,  545 
vapor  tension  of,  121 
Hydrazine,  422 
Hydrion,  properties  of,  346 
Hydrocarbons,  490 
properties  of,  493 
Hydrogele,  523 
Hydrogen,  92 

chemical  properties,  108 
physical  properties,  106 
preparation,  93,  95,  97,  99,  362 
bromide,  preparation,  230 

properties,  232 

chloride,  chemical  properties,  183 
composition,  183 
physical  properties,  182 
preparation,  178,  585 
fluoride,  physical  properties,  241 

preparation,  241 
iodide,  dissociation  of,  254 

interaction  with  sulphuric  acid,  237 
preparation,  237 
properties,  239 
pentasulphide,  376 
peroxide,  126,  303 
selenide,  401 
sulphate,  properties,  387 
sulphide,  371 

chemical  properties,  373,  374 
telluride,  403 

Hydrolysis,  definition  of,  181 
of  halogen  compounds,  534 
of  salts,  344 


774 


INDEX 


Hydrosole,  523 

Hydroxidion,  properties  of,  349 
Hydroxylamine,  423 
Hypobromites,  276 
Hypochlorites,  265 
Hypotheses,  formulative,  141 
Hypotheses,  stochastic,  142 
Hypothesis,  128 

atomic,  217 

Avogadro's,  131,  190 

corpuscular,  222 

ionic,  296,  317 

kinetic-molecular,  128 

molecular  applied  to  solutions,  150 

ICE,  115 

Iceland  spar,  590 
Identification,  means  of,  35,  39 
Illuminating-gas,  composition  of,  513 
Indicators,  353,  355 
Indigo,  688 
Indium,  681 

Ink,  754;  sympathetic,  759 
Insoluble  compounds,  theory  of  precip- 
itation of,  597,  601,  644 

theory  of  solution  of,  585,  597,  600, 

621,  644,  648 

Internal  rearrangement,  15,  488 
Inversion,  500 
Iodine,  233 
Iodine,  chlorides,  244 

pentoxide,  278,  279 

preparation,  234 

properties,  235 

uses,  236 
lodimetry,  236 
lodoform,  494 
Ion,  common,  effect  of  adding,  580 

definition  of,  321 

-product  constant,  583,  584 
Ionic  concentration,  579 

hypothesis,  296,  317 
objections  to,  319 
lonization,  degrees  of,  329 
lonogens,  296 
Iridium,  764 
Iron,  746 

alum,  756 

carbonyls,  757 

cast,  748,  751 


Iron,  chemical  properties,  752 

cyanides,  756 

wrought,  749 

see  Ferrous  and  Ferric 
Isomers,  262,  488,  696 
Isomorphism,  545 

KAINITE,  360,  559 
Kalion,  561 
Kaolin,  525,  689 
Kerosene,  492 
Kieserite,  645 
Krypton,  437 

LAKES,  689 
Lanthanum,  681,  682 
Lampblack,  476 
Law,  definition  of,  7 

Dulong  and  Petit's,  201 

Faraday's,  315 

Gay-Lussac's,  125 

Henry's,  154 

Le  Chatelier's,  260 

misconceptions    in     regard    to    the 
nature  of,  7 

of  combining  weights,  48 

of  mobile  equilibrium,  260 

of  partition,  155 

periodic,  410 

van  't  Hoff's,  260 
Lead,  698 

acetate,  704 

carbonate,  704 

chlorides,  700 

chromate,  727 

hydroxide,  701 

iodide,  700 

nitrate,  703 

oxides,  700,  701,  702 

sugar  of,  704 

sulphate,  704 

sulphide,  705 

Le  Blanc  soda  process,  572 
Light,  see  Photochemistry 
Lignin,  395 
Lime,  chloride  of,  see  Bleaching  powder 

milk  of,  595 

superphosphate  of,  467,  605 

-water,  68,  595 
Limestone,  590,  591,  594 


INDEX 


775 


Liquefaction  of  gases,  432 
Litharge,  700 
Lithium,  577 
Litmus,  70,  355 
Lodestone,  756 
Lunar  caustic,  631 

MAGNALIUM,  683 
Magnesia  alba,  644 
Magnesium,  642 

ammonium  arsenate,  645 

carbonate,  644 

chloride,  643 

hydroxide,  644 

oxide,  643 

phosphates,  645 

sulphate,  644 

sulphide,  645  , 
Malachite,  616 
Manganates,  737,  741 
Manganese,  737 

dioxide,  63,  171,  398,  739 

oxides  of,  738 

Manganic  compounds,  737,  740 
Manganites,  737,  741 
Manganous  compounds,  737,  739,  740 

sulphate,  739 
Marsh  gas,  493,  709,  714 
Mass  action,  250 
Massicot,  700 
Matches,  460 
Matrix,  367 
Matter,  23 

Mercur-ammonium  compounds,  659 
Mercuric,  cyanide,  658 

fulminate,  658 

oxide,  12,  61 

thiocyanate,  658 
Mercury,  652 

chlorides,  654 

iodides,  656 

nitrates,  657 

oxides,  656 

sulphides,  657 

Metallic  elements,  119,  337,  404,  534 
Metals,  electrical  conductivity  of,  532 

extraction  from  ores,  540 

hydroxides  of,  541 

occurrence  of,  539 

oxides  of,  541 


Metals,  physical  properties  of,  530 

see  Metallic  elements 
"Metals,"  recognition  of,  in  analysis, 

660 

Metastable  conditions,  159 
Methane,  493 
Methyl,  formate,  504 

orange,  355 
Mica,  525 

Microcosmic  salt,  467,  567  ^*" 
Migration,  ionic,  312 

speed  of,  314 
Minium,  701 

Mobile  equilibrium,  law  of,  260 
Mohr's  salt,  754 
Molar,  volume,  195 

weight,  194 
Mole,  194 
Molecular,  compounds,  443 

magnitudes,  140 

weight,  chemical  unit  of,  196 

weights,  193 

in  solution,  289,  292,  305 
of  elements,  204 
rule  for  measuring,  195 
Molecule,  the  chemical,  193,  296 
Molybdenum,  compounds  of,  732 
Mordants,  689 
Mortar,  596 
Multiple  proportions,  law  of,  41 

NAPHTHA,  492 

Nascent  state,  272,  423,  446 

Natrion,  577 

Neon,  437 

Neodymium,  682 

Nernst  lamp,  706 

Nessler's  reagent,  659 

Neutralization,  266,  351 

theory  of,  353 

thermochemistry  of,  357 

volume  change  in,  358 
Nickel,  700    760 

carbonyl,  762 

compounds,  761,  762 
Niobium,  721 
Nitrates,  442 
Nitric  oxide,  442 
Nitrites,  449 
Nitrogen,  62,  415 


776 


INDEX 


Nitrogen,  chloride,  424 

iodide,  425,  452 

oxides,  438  " 

oxygen  acids  of,  438 

tetroxide,  444 

Nitroglycerine,  441,  452,  504 
Nitrosyl  chloride,  448 
Nitrous  oxide,  450 
Nomenclature,  acids  and  salts,  263 

ionic  hypothesis,  296,  321 
Non-ionic  actions,  364 
Non-metallic  elements,  119,  405,  534 

OCTANTS,  law  of,  406 

Oleum,  388 

Orpiment,  712 

Osmium,  763 

Oxidation,  70,  72,  110,  454 

Oxides,  70 

Oxygen,  61 

atomic  weight  why  sixteen,  196 

chemical  properties,  67 

physical  properties,  66 

preparation,  63 
Ozone,  300 
Ozocerite,  492 

PALLADIUM,  764 

Paper,  manufacture  of,  395,  501,  686 

Paris  green,  624,  712 

Parke's  process,  627 

Pattinson's  process,  627 

Pearl  ash,  558 

Perfect  gas,  85 

Periodates,  278 

Periodic  system,  407 

applications  of,  412 
Peroxides,  308 
Petroleum,  491 

ether,  492 
Pewter,  694 
Phase  rule,  592 
Phases,  156 
Phenol,  441 
Phenolphthalein,  355 
Phlogiston,  9 
Phosgene,  487 
Phosphine,  461,  469 
Phosphonium  compounds,  462 
Phosphorus,  455 


Phosphorus,  acids  of,  465 

halides,  462 

oxides,  464 

pentachloride,  462 
dissociation  of,  255 

pentoxide,  464 

preparation,  455 

properties,  457 

red,  457 

sulphides,  471 

trichloride,  462 

trioxide,  464 

uses,  460 

yellow,  457 
Photochemistry,  458,  484,  633 

of  plants,  483 
Photography,  633 

Physics  needed  in  the  study  of  chem- 
istry, 39,  65 
Pink-salt,  695 
Plaster  of  Paris,  603v 
Platinum,  765 

compounds  of,  766 
Plumbates,  702 
Plumbic,  see  Lead 
Polarization,  324,  675 
Polymerization,  242 
Polysulphides,  376 
Porcelain,  690 
Potassamide,  420 
Potassium,  549 

aluminium  sulphate,  687 

arsenite,  712 

bromate,  556 

bromite,  552 

carbonate,  557 

chlorate,  64,  *275,  555.  583  ' 

chloride,  550 

chromate,  723 

cyanate,  507,  559 

cyanide,  558 

dichromate,  725 

ferrate,  758 

ferricyanide,  757 

ferrocyanide,  756 

fluorides,  552 

hydride,  550 

hydrogen  sulphate,  560 
sulphide,  560 
tart  rate,  561 


INDEX 


777 


Potassium,  hydroxide,  552,  558 

iodate,  556 

iodide,  551 

metantimoniate,  716 

nitrate,  556 

oxides,  555 

percarbonate.  558 

perchlorate,  556 

permanganate,  742 

plumbate,  702 

pyrosulphate,  560 

sulphate,  559 

sulphides,  560 

thiocyanate,  508,  559 

zincate,  648 
Potential,  chemical,  678 

differences  (anioris),  676 
(cations),  670 

discharging,  675 
Precipitate,  13 
Precipitation,    in    ionic    actions,    339, 

350 
'  rule  for,  597 

rule  of  (Berthollet's),  259 
Pressure,  dissociation,  257 

osmotic,  151,  283,  299 

partial,  88 

solution,  152 

vapor,  115,  135 

Principles,  summary  of,  188,  262 
Properties,  specific.  35 
Proustite,  626,  713 
Prussian  blue,  757 
Pumice-stone,  525 
Pure,  chemically,  definition  of>  34 
Putrefaction,  501 
Pyrargyrite,  626,  717 
Pyrite,  42,  756 
Pyrosulphates,  388 

QUARTATION,  532 

Quartz,  522 
Quicklime,  591,  594 
Quicksilver.  653 

RADICALS,  93 

organic,  494,  499 

valence  of,  104 
Radio-activity,  733 
Radium,  733 


Realgar,  712 
Red,  fire,  658 

heat,  temperature  corresponding  to, 
73 

lead,  701 

Reduction,  72,  110 
Reversible  actions,  176,  179,  237,  246 
Rhodium,  764 
Rhom-bohedron,  14 
Roasting,  378 
Rock  crystal,  522 
Rouge,  755 
Rubidium,  563 
Ruthenium,  763 

SALAMMONIAC,  564 
Saltpeter,  438,  569 
Salts,  264,  281,  337 

acid,  358 

basic,  359 

composition  of  hydrates  of,  545 

definition  of  term,  365 

double,  360 

general  methods  of  making,  542 

ionic  double  decompositions  of,  337 

mixed,  359 

nomenclature  of,  263 

of  complex  acids,  360,  363 

of  hydrogen,  345 

of  hydroxyl,  348 

precipitation  of,  339,  597,  601 

solubilities  of,  157,  544 
Samarium,  681 
Sand,  522 
Saponification,  505 
Scandium,  681 
Scheele's  green,  624,  712 
Schlippe's  salt,  717 
Sciences,  classification  of,  3 
Scientific  method,  4,  6,  10,  16,  17,  23, 

28-30,  128,  141,  191 
Selenium,  401 

Semi-permeable  partitions,  284,  299 
Schoenite,  559 
Silicates,  524 
Silicon,  518 

carbide,  520 

dioxide,  522 

hydride,  519 

tetrachloride,  520 


778 


INDEX 


Silicon,  tetrafluoride,  521 
Silver,  626 

arsenate,  632 

carbonate,  631 

chloride,  14 

chromate,  632,  728 

complex  compounds  of,  629 

cyanide,  630,  632 

halides,  629,  633 

nitrate,  13,  631 

orthophosphate,  632 

oxides  of,  630 

-plating,  632 

properties,  628 

sulphate,  632 

sulphide,  632 

thiosulphate,  630 

Simultaneous  actions,  231,  272,  275 
Slag,  540 
Soap,  505 
Soda,  crystals,  573 

washing,  573 
Sodium,  569 

aluminate,  99,  685 

amalgam,  569 

arsenite,  712 

carbonate,  manufacture,  571 
properties,  575 

chloride,  13,  570 

chromate,  724 

dichromate,  725 

hydride,  569 

hydrogen  carbonate,  575 

periodate,  279 

hydroxide,  570 

hyposulphite,  393 

"  hyposulphite,"  see  Thiosulphate 

nitrate,  14,  571 

nitrite,  571 

oxide,  571 

pentasulphide,  376 

peroxide,  303,  308,  571 

phosphates,  577 

plumbite,  701 

pyroantimoniate,  577,  716 

silicate,  577 

stannate,  696 

sulphate,  576 

hydrated,  120,  158,  160 

tetraborate,  528,  577 


Sodium,  thiosulphate,  576 

zincate,  99 
Soldering,  421 
Solubility,  curves  of,  157,  158 

data,  147,  154,  157,  531 

independent,  153 

influence  of  temperature  on,  156 

limits  of,  146 

measurement  of,  147 

of  gases,  153 

product,  585,  597 

separation  of  substances  by,  273 
Solution,  145 

definition  of,  165 

heat  of,  164 

in  two  solvents,  155 

of  insoluble  substances,  597,  598,  601, 
621,  644,  648 

scope  of  word,  146 
Solutions,  boiling-point  of,  162 

densities  of,  164 

freezing-points  of,  163 

general  properties  of,  145 

is-osmotic,  285 

molar,  149 

normal,  148 

saturated,  148,  153,  161,  581 

solid,  146 

standard,  236 

supersaturated,  159 

vapor  tension  of,  161 
Solution  tension,  hypothesis  of,  670 
Solvay  soda  process,  574 
Specific  heats  of  elements,  211 
Spectroscope,  561 
Spectrum,  562 
Speed  of  reaction,  111,  251,  394 

affected  by,  catalysis,  75 
concentration,  74,  250,  394 
solution,  76,  283 
temperature,  73 
Spinelles,  686 
Spirit  of  hartshorn,  418 
Stable,  meaning  of  term,  119 
Stannous,  see  Tin 
Starch,  500 
Steam,  115 
Stearin,  506 
Steel,  properties  of,  750 

theory  of  tempering,  750 


INDEX 


779 


Stibine,  714 
Stibnite,  714 
Storage  battery,  702 
Strontium,  chloride,  608 

hydrated,  121 
chromate,  728 
compounds,  608 
peroxide,  305 
Structure,    molecular,    204,    224,    see 

Constitution. 
Sublimation,  235,  463 
Substance,  meaning  of  term,  30,  32,  35 
Substitution,  176 
Suint,  557 

Sulphantimoniates,  717 
Sulphantimonites,  717 
Sulpha  rsenates,  712 
Sulpha  rsenites,  712 
Sulphates,  390 

acid,  390,  see  Hydrogen  sulphates 
Sulphides,  375 

relative  solubilities  of,  651 
Sulphites,  395 

Sulphocyanates,  see  Thiocyanates 
Sulphostannate,  697 
Sulphur,  11,367 

chemical  properties,  370 
chemical  relations,  377 
crystalline  form,  11 
dioxide,  378 

family,  chemical  relations  of,  403 
monochloride,  398 
oxides  of,  378 
oxygen  acids  of,  381  • 
physical  properties,  368 
tetrachloride,  399 
trioxide,  380 
Sulphuretted  hydrogen,  see  Hydrogen 

sulphide 

Sulphuryl  chloride,  399 
Summary  of  principles,  188,  262 
Superphosphate  of  lime,  467,  605 
Sylvite,  551 
Symbols,  53 
Synthesis,  473 
System,  meaning  of  term,  158 

TANTALUM,  721 
Tartar-emetic,  716 
Tellurium,  402 


Temperature,  critical,  134 
Tests,  99 
Thallium,  681 
Theory,  see  Hypothesis 
Thermochemistry,  76,  233,  271,  307, 357 
Thionyl  chloride,  399 
Thiocyanates,  250,  508,  559,  757 
Thorium,  510,  705 
Tin,  693 

bromide,  695 

chloride,  694,  695 

hydroxide,  696 

oxides,  696 

-stone,  693 

sulphides,  697 
Titanium,  705 
Titration,  352 
Touch-paper,  557 
Transformation  by  steps,  453 
Transition  points,  115,  118,  369,  593 
Triethylamine,  156 
Tungsten,  732 
TurnbulPs  blue 
Turquoise,  683 
Type-metal,  699 

ULTRAMARINE,  690 
Units,  25,  58,  533 

electrical,  665 
Uranium,  compounds,  733 
Urea,  487 

VALENCE,  101 

definitions  of,  102,  103,  106 

exceptional,   106 

multiple,  105 

of  radicals,  104 
Vanadium,  721 
Vapor,  densities,  90 

pressure,  115,  135 

tension  of  solutions,  161 
Velocity,  see  Speed  of  reaction 
Venetian  red,  755 
Verdigris,  624 
Vermilion,  658 
Vinegar,  498 
Vitriol,  blue,  120,  625 

green,  754 

oil  of,  see  Acid,  sulphuric 

white,  649 


780 


INDEX 


Vitriols,  546,  649 
Volume,  atomic,  407 

molar,  195 
Volumetric  analysis,  353 

WATER,  113 

chemical  properties,  118 

composition,  124 

electrolysis,  676 

gas,  485 

hard,  113,  593 

ionization  of,  331 

mineral,  113 

of  crystallization,  see  Hydrates 

physical  properties,  114 

purification,  113 
Weight,  molar,  194 

Weights,    atomic,    see    under    Atomic 
weights 

molecular,      see     under     Molecular 

weights 
Weldon  process,  741 


Welsbach  mantles,  705 
Whiskey,  477,  502 
White  lead,  704 
Wine,  501 

XENON,  437 

YTTERBIUM,  681 
Yttrium,  681 

ZINC,  646 

acetate,  649 

-blende,  646,  649 

carbonate,  649 

chloride,  647 

hydroxide,  648 

oxide,  646,  648 

sulphate,  649 

sulphide,  649 

-white,  648 
Zincates,  648 
Zirconium,  705 


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